Atoms

At the heart of modern chemistry lies an assumption, a model really, that all chemists take for granted. The assumption had its origin with the ancient Greeks, such as Democritus, who enjoyed speculating about the nature of the universe while sipping their wine at the dinner table. The word atoms comes from the Greek, atomos, meaning uncuttable. According to some Greek philosophers, the universe consists of tiny pieces of matter that themselves cannot be further subdivided. Thus, a piece of paper can be torn into many small pieces and with a sharp knife can be further divided into even smaller pieces, but at some point the cutting must stop. At this point, the pieces would be extremely small, and these were thought to be the building blocks from which all matter was built. Another school of thought, the Aristotelian group, believed that matter was infinitely divisible, that one could take the smallest piece of matter imaginable and then divide it into smaller pieces, and in turn divide those pieces, and on and on. Although the atomic school of thought is the one that dominates all of science, it may come as a surprise to you that this matter is probably by no means settled. The theoretical physicists have found numerous particles that are pieces of the atom, and seem to find more every year. [At the moment many physicists believe that one of these particles, the quark, is the fundamental particle from which all other matter is constructed.

The Greeks were not experimentalists; it was not until about the tenth century and the beginning of Alchemy that chemistry progressed as we think of it today; that is, as an experimental science. By the 18th century scientists had done thousands of experiments on the properties of matter and had formulated laws such as the Law of Conservation of Matter, and the Law of Multiple Proportions. Moreover, it was known that matter can be divided into mixtures and substances, with substances being either compounds or elements (see Figure 6). Compounds can be separated into their constituent elements, but elements cannot be further subdivided. Thus, atoms are the building blocks of matter on the atomic scale, elements are the building blocks of matter on the macroscopic level.

Figure 6. The classification of matter.

These laws allowed John Dalton, an English school teacher, to formulate a more modern version of the atomic model. Dalton was able to explain these and other Laws by making the following assumptions:

1. All matter consists of small, indivisible particles (atoms).
2. All of the atoms of the same element are the same; that is, have the same size, weight, color, etc. Atoms of different elements are different.
3. Compounds consist of combinations of atoms of different elements in whole number ratios.

For example, according to Dalton, common table salt (sodium chloride) consists of a one to one ratio of sodium atoms to chlorine atoms. In fact, Dalton probably visualized sodium chloride pretty much as shown in Figure 7.

Figure 7. Dalton's view of sodium chloride.

Dalton's atomic theory helped chemists to understand the laws that had been formulated and many of the observations that had been made. Like most theories, however, within 50 years it was shown to be wrong in several respects.

Around the turn of the century, chemists found evidence that atoms are not the solid, indivisible entities that Dalton imagined. The three particles that "live" inside the atom are: a) the electron, a tiny, negatively charged particle that controls how atoms bond to one another, b) the proton, a much heavier, positively charged particle, and c) the neutron, a particle of essentially the same mass as the proton, but without a charge. After these subatomic particles were discovered, chemists and physicists naturally speculated about the structure of the atom. How were the electrons and protons arranged? After all, electrons and protons are oppositely charged and should be strongly attracted to one another. One hypothesis was that the electrons and protons were evenly distributed throughout the atom. This was disproved by experiments by Rutherford in the early part of this century, in which he directed alpha particles (positively charged particles that had been discovered in naturally-occuring radioactivity) toward a thin metal sheet (see Figure 8). Some of the alpha particles were deflected back to the source in a way that could only be explained by a high concentration of positive charge in the center of the atom.

Figure 8. The Rutherford alpha particle scattering experiment.

Rutherford's hypothesis, then, was that all of the protons and neutrons reside in a very small center or nucleus of the atom and that the electrons are outside of the nucleus. If an atom were the size of a football field, the nucleus would be about the size of a baseball at the center of the field, and the electrons would be the size of grains of sand around the outside of the field. Clearly, according to this, the currently accepted model, most of the atom is empty space.

One of the most important consequences of the discovery of the subatomic particles was the realization that the atoms of different elements contain different numbers of protons. For example, all atoms of carbon contain six protons, all atoms of oxygen contain eight protons. The number of protons in the nucleus of an atom is called the atomic number. A portion of the periodic table of the elements is shown in Figure 9. Notice that in this table the atomic number is given above the symbol for each element; the elements are arranged in order of increasing atomic number from left to right and down the table.

1

18

1 H 1.0079

2

13

14

15

16

17

2 He 4.0026

3 Li 6.941

4 Be 9.0122

5 B 10.811

6 C 12.011

7 N 14.007

8 O 15.999

9 F 18.998

10 Ne 20.180

11 Na 22.990

12 Mg 24.305

3

4

5

6

7

8

9

10

11

12

13 Al 26.982

14 Si 28.086

15 P 30.974

16 S 32.065

17 Cl 35.453

18 Ar 39.948

19 K 39.098

20 Ca 40.078

21 Sc 44.956

22 Ti 47.867

23 V 50.942

24 Cr 51.996

25 Mn 54.938

26 Fe 55.845

27 Co 58.933

28 Ni 58.693

29 Cu 63.546

30 Zn 65.409

31 Ga 69.723

32 Ge 72.64

33 As 74.922

34 Se 78.96

35 Br 79.904

36 Kr 83.798

37 Rb 85.468

38 Sr 87.62

39 Y 88.906

40 Zr 91.224

41 Nb 92.906

42 Mo 95.94

43 Tc (98)

44 Ru 101.07

45 Rh 102.91

46 Pd 106.42

47 Ag 107.87

48 Cd 112.41

49 In 114.82

50 Sn 118.71

51 Sb 121.76

52 Te 127.60

53 I 126.90

54 Xe 131.29

55 Cs 132.91

56 Ba 137.33

57 - 71

72 Hf 178.49

73 Ta 180.95

74 W 183.84

75 Re 186.21

76 Os 190.23

77 Ir 192.22

78 Pt 195.08

79 Au 196.97

80 Hg 200.59

81 Tl 204.38

82 Pb 207.2

83 Bi 208.98

84 Po (209)

85 At (210)

86 Rn (222)

87 Fr (223)

88 Ra (226)

89 - 103

104 Rf (261)

105 Db (262)

106 Sg (266)

107 Bh (264)

108 Hs (277)

109 Mt (268)

110 Ds (271)

111 Rg (272)

57 La 138.91

58 Ce 140.12

59 Pr 140.91

60 Nd 144.24

61 Pm (145)

62 Sm 150.36

63 Eu 151.96

64 Gd 157.25

65 Tb 158.93

66 Dy 162.50

67 Ho 164.93

68 Er 167.26

69 Tm 168.93

70 Yb 173.04

71 Lu 174.97

89 Ac (227)

90 Th 232.04

91 Pa 231.04

92 U 238.03

93 Np (237)

94 Pu (244)

95 Am (243)

96 Cm (247)

97 Bk (247)

98 Cf (251)

99 Es (252)

100 Fm (257)

101 Md (258)

102 No (259)

103 Lr (262)

Figure 9. The periodic table of the elements.

Problem One

How many protons are there in the nucleus of an atom of the element calcium (Ca)?

press for answer 

With the invention of instrumentation such as the mass spectrometer, which allows atoms of different mass to be separated and identified, another discovery was made: atoms of the same element can have different numbers of neutrons. For example, there are two naturally-occurring carbon atoms: one has 6 neutrons, the other has 7 neutrons. The total number of protons and neutrons is called the mass number. Two atoms of the same element that contain different numbers of neutrons are called isotopes. The mass number of an element is frequently designated as a superscript on the left side of the symbol for the element. Thus, ¹³C is the isotope of carbon that contains 7 neutrons. Sometimes, the atomic number and the mass number are designated at the same time, the atomic number as a subscript and the mass number as a superscript.

The number of electrons in an atom is determined by the charge on the atom. If the atom is neutral, that is, has no charge, then the positive charge produced by the protons must be balanced by the negative charge of the electrons. Thus, a neutral carbon atom must contain six electrons. However, if the carbon atom has a positive charge of one it will contain one less electron. Positively charged atoms have more protons than electrons and are called cations. Negatively charged atoms have more electrons than protons and are called anions. The charge of an atom is designated by a superscript on the right hand side of the symbol of the element. A carbon atom that has eight electrons would be designated as C^2-.

Problem Two

What is the symbol for the ion with an atomic number of 12 that has 10 electrons?

press for answer 

In addition to the atomic number, the atomic weight of each element is also given, located below the symbol for the element on the periodic chart. Generally, chemists measure the mass of things in grams or kilograms (a kilogram is a thousand grams, a milligram is a thousandth of a gram). In fact, almost everything that a scientist measures has units; if speed is being measured, the units may be meters per second (m/s); if volume is measured it is expressed as cubic meters, or cubic centimeters (a cubic centimeter is a milliliter). Without these units the scientific community would have no idea of what the number designated. Suppose for example, that a chemist reports that 32 of sulfur is mixed with 23 of magnesium. What is the chemist referring to? Volume--if so does she mean 32 milliliters (mL) or 32 liters? Or is it mass--if so, is it 32 grams (g) of sulfur or 32 milligrams? It could even be 32 pounds of sulfur if the chemist does not use the metric system. Because these units are so important, scientists have agreed to use the metric system of measurement and prefixes for the metric units that are shown in Table 1. In certain applications, non metric units may still be used, however. In 1960, the General Conference of Weights and Measurements recommended that a single unit be used for each measured quantity. This Systeme International d' Unites (International System of Units, SI) is based on the metric system and consists of the seven fundamental base units shown in Table 2.

Table 1. Metric Prefixes

Factor	Prefix	Symbol
109	giga	G
106	mega	M
103	kilo	k
10-1	deci	d
10-2	centi	c
10-3	milli	m
10-6	micro	μ
10-9	nano	n
10-12	pico	p
10-15	femto	f

Table 2. SI Units

Quantity	Name of Unit	Symbol
Length	meter	m
Mass	kilogram	kg
Time	second	s
Electric current	ampere	A
Thermodynamic temperature	kelvin	K
Luminous intensity	candela	cd
Amount of substance	mole	mol

  For example, 1 MeV = 10⁶ electron volts; 1 pm = 10^-12 meters; 1 mg = 10^-3 grams.

Other units can be derived from the SI base units. For example, the SI unit for volume is cubic meters (m³); the SI unit for force is the newton, N, which is obtained when mass is expressed in kilograms, and acceleration is given in meters per second; pressure is given in pascals (Pa), which is a newton per square meter. The complete adoption of SI units has not yet occurred. Pressure, for example, is usually expressed as torr, millimeters of mercury, or atmospheres (in some cases even in the English units of pounds per square inch, psi). Likewise, volume is usually expressed in liters (L) which is 1000 cm³; temperature is often given in degrees Celsius ( or centigrade, °C) or even in the English units of Fahrenheit (°F).

At any rate, the atomic weight given below the element in the periodic chart does not contain a unit. Of course, this could simply be a result of a decision on the part of the illustrator, an assumption, perhaps, that everyone would know that the weight was given in units of, say, tons. In fact, these weights really do not have units, they are simply relative. Hundreds of chemists worked for many years to establish these relative atomic weights. Originally, the weights were adjusted so that hydrogen, the lightest element, would have a weight of 1. If you compare hydrogen and oxygen, you will notice that an atom of oxygen is 16 times heavier than an atom of hydrogen. This arbitrary scale was adjusted later on, after the invention of the mass spectrometer and the discovery of isotopes. The weight of the ¹²C nucleus was then assigned a value of exactly 12 (12.00000). Notice that because there is a little ¹³C present in carbon, the atomic weight of carbon is not 12.000000, but 12.01.

As you might imagine someone started to wonder about the absolute mass of atoms. How much does one ¹²C atom weigh? This weight can be expressed as 12.0000 atomic mass units (amu), but unless we know how much an atomic mass unit is, this is not particularly helpful. This problem was solved by determining the number of atoms of an element in an amount of the element equal to its atomic weight in grams. This number, called Avogadro's number in honor of the famous Italian scientist, Amadeo Avogadro, is huge--602200000000000000000000. Because numbers such as this are very cumbersome to read they are usually expressed in scientific notation. This number is given as 6.022 x 10²³, which means that 6.022 is multiplied by 10 raised to the 23^rd power; in other words, that 6.022 is multiplied by a 1 followed by 23 zeros. The magnitude of this number can be appreciated by imagining a gigantic wall constructed of Avogadro's number of grains of sand (each 0.05 cm in diameter). If this wall were one mile wide and one mile high, it would extend all the way around the earth (about 20,000 miles).

Problem Three

If you are not familiar with scientific notations, sometimes also called exponential notation, take another look at 6 x 10²³. It literally means that you multiply 6 by a 1 with 23 zeros to the right of it. If a number is expressed with a negative exponent; for example, 3.1 x 10^-3, means that 3.1 is multiplied by 1 with 3 places between it and the decimal place -- 0.001. When you multiply 3.1 by 0.001 you get 0.0031. Thus, 3.1 x 10^-3 is the same as 0.0031. Can you express 2.23 x 10² and 6.02 x 10^-5 in their decimal equivalents?

press for answer 

Phosphorus has an atomic weight of 31.0. In this 31.0 grams there are 6 x 10²³atoms. Clearly, one average atom of phoshorus must have a very tiny mass. Indeed, we can calculate this mass by dividing the mass of 6.02 x 10²³ atoms (31.0 g) by the number of atoms:

(31.0 g)/(6 x 10²³) = 5 x 10^-23 g

Thus, one atom of phosphorus weighs 5 x 10^-23 g, a very tiny mass indeed.

Avogadro's number of atoms is so frequently used that it has been given a special name-- a mole. This term is similar to the use of dozen for 12 items or ream for 500 sheets of paper. Thus, in 31 grams of phosphorus there is a mole of phosphorus atoms (6 x 10²³ atoms). Five moles of carbon is 5 x 12 g = 60 grams of carbon. This 60 grams of carbon contains 5 x (6 x 10²³) = 30 x 10²³ atoms. Chemists are so accustomed to working with moles that you will hear them say things like: "The reaction was run on a 0.1 mole scale", or "We'll need at least a millimole in order to obtain a good spectrum."

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The Scientific Method

It is important to know something about the way in which new theories are produced. To illustrate this process, sometimes called the scientific method, let us follow the thoughts and actions of a scientist who happens upon the box pictured in Figure 4. The box is apparently constructed from some heavy metal and has three rather stiff ropes protruding from the holes labeled A, B, and C. Her curiosity aroused, our scientist attempts to open the box to discover its purpose and inner workings. She soon discovers, however, that she lacks the necessary tools and must be satisfied with observations made from outside the box.

Figure 4. A curious box.

An initial tug on rope C produces no apparent movement of ropes A and B. This observation triggers the thought that the box may contain three unconnected, independent ropes. This mental visualization is our scientist's first hypothesis of the nature of the interior of the box.

Now, she reasons, if this hypothesis is correct, each rope should move to its limit without affecting (moving) any of the other ropes. In order to test her hypothesis, she pulls each rope and observes the effect of that action on the other ropes. Rope C appears to move independently of the other, but the effects of ropes A and B on each other can be expressed as a mathematical law. If x represents the distance traversed by rope A, and y the distance traversed by rope B, the law becomes

x = 2y

Because of the specific interdependency of ropes A and B, the scientist now proposes the model pictured below in Figure 5. In this model, rope A is wound around a drum attached to an axle; rope B is wound around the axle itself. The circumference of the drum is greater than the circumference of the axle (in fact, 2 times as great), and rope A is therefore played out at a greater rate than rope B. Rope C remains unattached and independent.

Figure 5. A model of the interior of the box.

This model provides an explanation for all of the experimental data, and it also permits the formulation of new questions and predictions. For example: Can rope C be withdrawn from the box completely, or is it held inside by a knot at the end? Are the ends of ropes A and B attached to the drum and the axle? If the model is correct, then when rope B is withdrawn to its limit, rope A may not have reached its limit, and (assuming that the ropes are attached) a further pull on rope A may wind rope B back into the box. These questions suggest additional experiments which might never have been conceived without the help of the model.

While our scientist has spent a considerable amount of energy investigating an almost trivial problem, her approach to the problem contains many of the features of the "scientific method": experimental observation that leads to the formulation of a law, a hypothesis that leads to a model or theory, and the subsequent use of the model to design new experiments. Each hypothesis leads to a model, which may be discarded after additional experiments are performed, or, if the experiments are all consistent with the model, the model is retained until contradictory evidence is obtained.

Because of the central role of models, it is important to be cognizant of a number of their characteristics. First, the scientist usually draws on her own experiences in fashioning a theory. In our example, it might be suggested that the ropes are controlled by elves residing in the box, but our scientist has never seen an elf, nor does she believe in the existence of such creatures. On the other hand, she has seen mechanical devices such as winches that employ ropes on drums, and she has seen a spool of thread. Many models, designed to account for the behavior of matter so small that it has never been seen, are based on the behavior of macroscopic bodies, such as billiard balls, which lie within the realm of everyone's experience.

On the other hand, some models are mathematical and abstract in nature. For example, the mathematical nature of the contemporary model of the electron makes many of its features difficult to visualize. Indeed, some scientists feel that the most significant scientific discoveries occur within the realm of mathematics.

It is also important to realize that a given set of experiments and observations can usually be explained by more than one model. Our scientist could have developed a model based on gears rather than drums, and in fact there are a number of alternate models that will satisfactorily account for the behavior of the ropes. As data and observations accumulate, one of a set of equally good models may become more satisfactory than the rest, or the choice of model may be based on considerations of simplicity or symmetry or usefulness.

Finally, the fact that models may not, and very likely do not, correspond to reality cannot be overemphasized. Since the box cannot be opened, the scientist will probably never know if the box really does contain a drum and axle. When the model is intended as a picture or visualization of matter at the sub microscopic, molecular level, the problem is even more acute. Atoms cannot possibly be either billiard balls or mathematical abstractions, nor is it likely that atoms behave like billiard balls. And yet, the billiard ball model of atoms is at the heart of the determination of the structure of the nucleic acids DNA and RNA, the revelation of the genetic code, and all of its biological implications. Thus, while the correspondence between the model and reality may not be very high, the benefits of the model, the development of new experiments, the discovery of new laws of nature, and so forth - may be very great indeed.

Reference: C. Yoder, O. Retterer, M. Thomsen, and K. Hess, Interactive Chemistry, Mosby Year-Book, 1997.

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What is Chemistry?

Chemistry is frequently defined as the study of matter and the reactions that matter undergoes. Actually, physicists, geologists, and biologists also study matter, but only chemists study the reactions that matter undergoes. For example, only chemists make compounds and try to understand the reactions that produce the compounds. Indeed, a very large segment of chemists are employed by the chemical and pharmaceutical industry for the very purpose of preparing new plastics, coatings, ceramics, drugs, fillers, alloys, and so on. These synthetic chemists must first determine what reaction can be used to synthesize their target compound and then determine what conditions will optimize the yield of the compound in order to make the compound in the most cost-effective way. After the best reaction conditions have been determined, the chemist must determine how to purify the compound, and, finally, the chemist must identify it. This final process of identification usually includes not only being certain that the compound contains the right percentage of the various elements from which it is composed, but also involves the determination of the 3-dimensional structure of the compound.

Structural details are often crucial to the activity of the compound. For example, the compounds dextrophane and levorphane differ in a very subtle way. They are non-superimposable mirror images of one another in the same way that our hands are non-superimposable mirror images of one another. Yet, because of a quirk of the evolutionary process, our bodies are able to recognize this subtle difference and produce a very different response to the two compounds: levorphan is more strongly analgesic and addictive than morphine, whereas dextrophan is neither addictive nor an analgesic. Figure 1 shows the structural formulas of the two compounds (we will discuss the various types of formulas in a later section). In Figure 2, the same molecules are shown as computer generated molecular models. In part (a), ball and stick models are shown, while part (b) shows space-filling models.

Figure 1. Structural formulas for levorphan and dextrophan.

(a)

(b)

Figure 2. Molecular models of levorphan and dextrophan. (a) ball and stick models, (b) space-filling models.

When you compare Figure 1 and Figure 2 you will find that the lines in the structural formulas indicate the attachment of atoms to one another. These lines are called bonds. Some bonds are single bonds, some are double bonds, others are triple bonds. Generally the greater the number of bonds between two atoms, the stronger the attachment of the two atoms. Notice also that in the molecular model, the atoms have different colors and sizes. The colors are obviously used to distinguish one type of atom from another. Also, recognize that the distance of the atom from the reader (depth) is indicated by the size of the atom. The space-filling model is designed to give a somewhat more accurate representation of the molecule by portraying the space filled by the electrons around the atoms. Although these models are probably more realistic representations of the molecules, they are also more difficult to "read." Most chemists prefer to see the ball and stick models, but they use space-filling representations when they are interested in the spatial requirements of certain parts of a molecule.

Molecular modeling has become an important part of the arsenal of the synthetic chemist as well as the theoretical chemist. Frequently, the synthetic chemist makes use of computer modeling to identify compounds that will have certain physical properties or produce certain physiological responses.

Theoretical and physical chemists are concerned with the description of the bonding between atoms and understanding the changes in electronic structure that occur when a reaction takes place. They produce theories or models that are eventually incorporated into the body of chemistry and used by synthetic chemists to make compounds with new, and frequently useful, properties.

Chemists can also be categorized according to the traditional sub disciplines: inorganic (elements other than carbon), organic (carbon compounds), analytical (methods used to separate and identify compounds), and physical chemists. Today, there are also many other cross disciplinary areas that occupy chemists: biochemists try to understand and apply the chemistry of biological processes, materials chemists attempt to synthesize new materials such as superconductors or artificial skin, environmental chemists study the chemistry of the environment and monitor and solve environmental problems, while forensic chemists apply chemistry to the solution of crimes.

Figure 3. The disciplines of chemistry.

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An Introduction to Chemistry

This book is designed as a brief introduction to chemistry. Our primary objective is to give you enough understanding of the fundamental concepts and language of chemistry to allow you to read and understand articles written in newspapers such as the New York Times or in magazines such as Scientific American. Of particular importance is the concept of structure. An understanding of the shapes and motions of molecules and ions and the structure of matter in bulk is crucial to the form of the world in which we live. It is also important to know something of the language of chemistry: the meaning of words such as isotope, isomer, structural formula, the mole, the rate and extent of a reaction, equilibrium, and energy.

The approach to the topics will be informal and, we hope, not intimidating. There will be, however, the occasional problem to solve. We will also provide numerous illustrations and photographs to provide you with examples of different ways to view concepts.

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General Chemistry/Introduction

Chemistry is Everywhere

Chemistry: the study of the properties, composition, and transformation of matter.

The modern human experience places a large emphasis upon the material world. From the day of our birth to the day we die, we are frequently preoccupied with the world around us. Whether struggling to feed ourselves, occupying ourselves with modern inventions, interacting with other people or animals, or simply meditating on the air we breathe, our attention is focused on different aspects of the material world. In fact only a handful of disciplines—certain subsets of religion, philosophy, and abstract math—can be considered completely unrelated to the material world. Everything else is somehow related to chemistry, the scientific discipline which studies the properties, composition, and transformation of matter.

Branches of Chemistry

Chemistry itself has a number of branches:

• Analytical chemistry seeks to determine the composition of substances.
• Biochemistry is the study of chemicals found in living things (such as DNA and proteins).
• Inorganic Chemistry studies substances that do not contain carbon.
• Organic chemistry studies carbon-based substances. Carbon, as described in more detail in this book, has unique properties that allow it to make complex chemicals, including those of living organisms. An entire field of chemistry is devoted to substances with this element.
• Physical chemistry is the study of the physical properties of chemicals, which are characteristics that can be measured without changing the composition of the substance.

This is the structure of table salt, or sodium chloride.

Chemistry as a discipline is based on a number of other fields. Because it is a measurement-based science, math plays an important role in its study and usage. A proficiency in high-school level algebra should be all that is needed in this text, and can be obtained from a number of sources. Chemistry itself is determined by the rules and principles of physics. Basic principles from physics may be introduced in this text when necessary.

Why Study Chemistry?

There are many reasons to study chemistry. It is one pillar of the natural sciences necessary for detailed studies in the physical sciences or engineering. The principles of biology and psychology are rooted in the biochemistry of the animal world, in ways that are only now beginning to be understood. Modern medicine is firmly rooted in the chemical nature of the human body. Even students without long-term aspirations in science find beauty in the infinite possibilities that originate from the small set of rules found in chemistry.

Chemistry has the power to explain everything in this world, from the ordinary to the bizarre. Why does iron rust? What makes propane such an efficient, clean burning fuel? How can soot and diamond be so different in appearance, yet so similar chemically? Chemistry has the answer to these questions, and so many more. Understanding chemistry is the key to understanding the world as we know it.

This Book: General Chemistry

An introduction to the chemical world is set forth in this text. The units of study are organized as follows.

1. Properties of Matter: An explanation of the most fundamental concept in chemistry: matter.
2. Atomic Structure: While technically in the domain of physics, atoms determine the behavior of matter, making them a necessary starting point for any discussion of chemistry.
3. Compounds and Bonding: Chemical bonding is introduced, which explains how less than one hundred naturally-occurring elements can combine to form all the different compounds that fill our world.
4. Chemical Reactions: Things get interesting once chemical reactions start making and breaking bonds.
5. Aqueous Solutions: Substances dissolved in water have special properties. This is when acids and bases are introduced.
6. Phases of Matter: A detailed look at the organization of substances, with particular focus on gases.
7. Chemical Equilibria: Chemical reactions don't go on forever. Equilibrium is the balance that reactions seek to achieve.
8. Chemical Kinetics: Kinetics explain why it takes years for an iron nail to rust, but only a split second for a hydrogen-filled hot air balloon to explode.
9. Thermodynamics: Two things decide which reactions can occur and which reactions cannot: heat and chaos. Or enthalpy and entropy, as they are called in thermodynamics
10. Chemistries of Various Elements: An exploration of the elements that make up all substance. Includes an introduction to nuclear chemistry and carbon, the essence of organic chemistry.

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