Chem Honors Unit 1

2.1 Atomic Theory

In the fifth century B.C. the Greek philosopher Democritus expressed the belief that all matter consists of very small, indivisible particles, which he named atomos (meaning uncuttable or indivisible). Although Democritus’ idea was not accepted by many of his contemporaries (notably Plato and Aristotle), somehow it endured. Experimental evidence from early scientific investigations provided support for the notion of “atomism” and gradually gave rise to the modern definitions of elements and compounds. It was in 1808 that an English scientist and schoolteacher, John Dalton, formulated a precise definition of the indivisible building blocks of matter that we call atoms. Dalton’s work marked the beginning of the modern era of chemistry. The hypotheses about the nature of matter on which Dalton’s atomic theory is based can be summarized as

Elements are composed of extremely small particles, called atoms.
All atoms of a given element are identical, having the same size, mass, and chemical properties. The atoms of one element are different from the atoms of all other elements.
Compounds are composed of atoms of more than one element. In any compound, the ratio of the numbers of atoms of any two of the elements present is either an integer or a simple fraction.
A chemical reaction involves only the separation, combination, or rearrangement of atoms; it does not result in their creation or destruction. Figure 2.1 is a schematic representation of hypotheses 2 and 3. Dalton’s concept of an atom was far more detailed and specific than Democritus’. The second hypothesis states that atoms of one element are different from atoms of all other elements. Dalton made no attempt to describe the structure or composition of atoms—he had no idea what an atom is really like. But he did realize that the different properties shown by elements such as hydrogen and oxygen can be explained by assuming that hydrogen atoms are not the same as oxygen atoms. The third hypothesis suggests that, to form a certain compound, we need not only atoms of the right kinds of elements, but the specific numbers of these atoms as well. This idea is an extension of a law published in 1799 by Joseph Proust, a French chemist. Proust’s law of definite proportions states that different samples of the same compound always contain its constituent elements in the same proportion by mass. Thus, if we were to analyze samples of carbon dioxide gas obtained from different Figure 2.1 (a) According to Dalton’s atomic theory, atoms of the same element are identical, but atoms of one element are different from atoms of other elements. (b) Compound formed from atoms of elements X and Y. In this case, the ratio of the atoms of element X to the atoms of element Y is 2:1. (b) Atoms of element X Atoms of element Y Compound of elements X and Y (a)

sources, we would find in each sample the same ratio by mass of carbon to oxygen. It stands to reason, then, that if the ratio of the masses of different elements in a given compound is fixed, the ratio of the atoms of these elements in the compound also must be constant. Dalton’s third hypothesis also supports another important law, the law of multiple proportions. According to this law, if two elements can combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in ratios of small whole numbers. Dalton’s theory explains the law of multiple proportions quite simply: The compounds differ in the number of atoms of each kind that combine. For example, carbon forms two stable compounds with oxygen, namely, carbon monoxide and carbon dioxide. Modern measurement techniques indicate that one atom of carbon combines with one atom of oxygen in carbon monoxide and that one atom of carbon combines with two oxygen atoms in carbon dioxide. Thus, the ratio of oxygen in carbon monoxide to oxygen in carbon dioxide is 1:2. This result is consistent with the law of multiple proportions because the mass of an element in a compound is proportional to the number of atoms of the element present (Figure 2.2). Dalton’s fourth hypothesis is another way of stating the law of conservation of mass, which is that matter can be neither created nor destroyed.† Because matter is made of atoms that are unchanged in a chemical reaction, it follows that mass must be conserved as well. Dalton’s brilliant insight into the nature of matter was the main stimulus for the rapid progress of chemistry during the nineteenth century.

2.2 The Structure of the Atom

On the basis of Dalton’s atomic theory, we can define an atom as the basic unit of an element that can enter into chemical combination. Dalton imagined an atom that was both extremely small and indivisible. However, a series of investigations that began in the 1850s and extended into the twentieth century clearly demonstrated that atoms actually possess internal structure; that is, they are made up of even smaller particles, which are called subatomic particles. This research led to the discovery of three such particles—electrons, protons, and neutrons. The Electron In the 1890s many scientists became caught up in the study of radiation, the emission and transmission of energy through space in the form of waves. Information gained from this research contributed greatly to our understanding of atomic structure. One device used to investigate this phenomenon was a cathode ray tube, the forerunner of the television tube (Figure 2.3). It is a glass tube from which most of the air has been evacuated. When the two metal plates are connected to a high-voltage source, the negatively charged plate, called the cathode, emits an invisible ray. The cathode ray is drawn to the positively charged plate, called the anode, where it passes through a hole and continues traveling to the other end of the tube. When the ray strikes the specially coated surface, it produces a strong fluorescence, or bright light.

† According to Albert Einstein, mass and energy are alternate aspects of a single entity called mass-energy. Chemical reactions usually involve a gain or loss of heat and other forms of energy. Thus, when energy is lost in a reaction, for example, mass is also lost. Except for nuclear reactions (see Chapter 21), however, changes of mass in chemical reactions are too small to detect. Therefore, for all practical purposes mass is conserved. Animation: Cathode Ray Tube ARIS, Animations cha4 2.2 The Structure of the Atom 31 In some experiments, two electrically charged plates and a magnet were added to the outside of the cathode ray tube (see Figure 2.3). When the magnetic field is on and the electric field is off, the cathode ray strikes point A. When only the electric field is on, the ray strikes point C. When both the magnetic and the electric fields are off or when they are both on but balanced so that they cancel each other’s influence, the ray strikes point B. According to electromagnetic theory, a moving charged body behaves like a magnet and can interact with electric and magnetic fields through which it passes. Because the cathode ray is attracted by the plate bearing positive charges and repelled by the plate bearing negative charges, it must consist of negatively charged particles. We know these negatively charged particles as electrons. Figure 2.4 shows the effect of a bar magnet on the cathode ray. An English physicist, J. J. Thomson, used a cathode ray tube and his knowledge of electromagnetic theory to determine the ratio of electric charge to the mass of an individual electron. The number he came up with is 1.76 108 C/g, where C stands for coulomb, which is the unit of electric charge. Thereafter, in a series of experiments carried out between 1908 and 1917, R. A. Millikan, an American physicist, found the charge of an electron to be 1.6022 1019 C. From these data he calculated the mass of an electron: which is an exceedingly small mass. Radioactivity In 1895, the German physicist Wilhelm Röntgen noticed that cathode rays caused glass and metals to emit very unusual rays. This highly energetic radiation penetrated matter, darkened covered photographic plates, and caused a variety of substances .9.10 X 1028 g mass of an electron. A cathode ray tube with an electric field perpendicular to the direction of the cathode rays and an external magnetic field. The symbols N and S denote the north and south poles of the magnet. The cathode rays will strike the end of the tube at A in the presence of a magnetic field, at C in the presence of an electric field, and at B when there are no external fields present or when the effects of the electric field and magnetic field cancel each other. Electrons are normally associated with atoms. However, they can also be studied individually. Animation: Millikan Oil Drop ARIS, Animations cha48518_c fluoresce. Because these rays could not be deflected by a magnet, they could not contain charged particles as cathode rays do. Röntgen called them X rays. Not long after Röntgen’s discovery, Antoine Becquerel, a professor of physics in Paris, began to study fluorescent properties of substances. Purely by accident, he found that exposing thickly wrapped photographic plates to a certain uranium compound caused them to darken, even without the stimulation of cathode rays. Like X rays, the rays from the uranium compound were highly energetic and could not be deflected by a magnet, but they differed from X rays because they were generated spontaneously. One of Becquerel’s students, Marie Curie, suggested the name radioactivity to describe this spontaneous emission of particles and/or radiation. Consequently, any element that spontaneously emits radiation is said to be radioactive. Further investigation revealed that three types of rays are produced by the decay, or breakdown, of radioactive substances such as uranium. Two of the three kinds are deflected by oppositely charged metal plates (Figure 2.5). Alpha () rays consist of 2 Atoms, Molecules, and Ions Figure 2.4 (a) A cathode ray produced in a discharge tube traveling from the cathode (left) to the anode (right). The ray itself is invisible, but the fluorescence of a zinc sulfide coating on the glass causes it to appear green. (b) The cathode ray is bent downward when the north pole of the bar magnet is brought toward it. (c) When the polarity of the magnet is reversed, the ray bends in the opposite direction. (a) (b) (c) – + α γ β Radioactive substance Lead block Figure 2.5 Three types of rays emitted by radioactive elements. rays consist of negatively charged particles (electrons) and are therefore attracted by the positively charged plate. The opposite holds true for rays—they are positively charged and are drawn to the negatively charged plate. Because rays have no charges, their path is unaffected by an external electric field. Animation: Alpha, Beta, and Gamma Rays ARIS, Animations cha 2.2 The Structure of the Atom 33 positively charged particles, called particles, and therefore are deflected by the positively charged plate. Beta () rays, or particles, are electrons and are deflected by the negatively charged plate. The third type of radioactive radiation consists of highenergy rays called gamma () rays. Like X rays, rays have no charge and are not affected by an external electric or magnetic field. The Proton and the Nucleus By the early 1900s, two features of atoms had become clear: They contain electrons, and they are electrically neutral. To maintain electrical neutrality, an atom must contain an equal number of positive and negative charges. On the basis of this information, Thomson proposed that an atom could be thought of as a uniform, positive sphere of matter in which electrons are embedded (Figure 2.6). Thomson’s so-called “plumpudding” model was the accepted theory for a number of years. In 1910 the New Zealand physicist Ernest Rutherford, who had earlier studied with Thomson at Cambridge University, decided to use particles to probe the structure of atoms. Together with his associate Hans Geiger and an undergraduate named Ernest Marsden, Rutherford carried out a series of experiments using very thin foils of gold and other metals as targets for particles from a radioactive source (Figure 2.7). They observed that the majority of particles penetrated the foil either undeflected or with only a slight deflection. They also noticed that every now and then an particle was scattered (or deflected) at a large angle. In some instances, an particle actually bounced back in the direction from which it had come! This was a most surprising finding, for in Thomson’s model the positive charge of the atom was so diffuse (spread out) that the positive particles were expected to pass through with very little deflection. To quote Rutherford’s initial reaction when told of this discovery: “It was as incredible as if you had fired a 15-inch shell at a piece of tissue paper and it came back and hit you.” To explain the results of the -scattering experiment, Rutherford devised a new model of atomic structure, suggesting that most of the atom must be empty space. This structure would allow most of the particles to pass through the gold foil with little or no deflection. The atom’s positive charges, Rutherford proposed, are all concentrated in the nucleus, a dense central core within the atom. Whenever an particle came close to a nucleus in the scattering experiment, it experienced a large repulsive force and therefore a large deflection. Moreover, an particle traveling directly toward a nucleus would experience an enormous repulsion that could completely reverse the direction of the moving particle. The positively charged particles in the nucleus are called protons. In separate experiments, it was found that the charge of each proton has the same magnitude as that of an electron and that the mass of the proton is 1.67262 1024 g—about 1840 times the mass of the oppositely charged electron. Positive charge spread over the entire sphere – – – – – – – Figure 2.6 Thomson’s model of the atom, sometimes described as the “plum-pudding” model, after a traditional English dessert containing raisins. The electrons are embedded in a uniform, positively charged sphere. Animation: -Particle Scattering ARIS, Animations Slit Detecting screen Gold foil (a) (b) α Particle emitter Figure 2.7 (a) Rutherford’s experimental design for measuring the scattering of particles by a piece of gold foil. Most of the particles passed through the gold foil with little or no deflection. A few were deflected at wide angles. Occasionally an particle was turned back. (b) Magnified view of particles passing through and being deflected by nuclei. cha48518_ch02_028-0 At this stage of investigation, scientists perceived the atom as follows. The mass of a nucleus constitutes most of the mass of the entire atom, but the nucleus occupies only about of the volume of the atom. We express atomic (and molecular) dimensions in terms of the SI unit called the picometer (pm), and A typical atomic radius is about 100 pm, whereas the radius of an atomic nucleus is only about 5 103 pm. You can appreciate the relative sizes of an atom and its nucleus by imagining that if an atom were the size of a sports stadium, the volume of its nucleus would be comparable to that of a small marble. Although the protons are confined to the nucleus of the atom, the electrons are conceived of as being spread out about the nucleus at some distance from it. The Neutron Rutherford’s model of atomic structure left one major problem unsolved. It was known that hydrogen, the simplest atom, contains only one proton and that the helium atom contains two protons. Therefore, the ratio of the mass of a helium atom to that of a hydrogen atom should be 2:1. (Because electrons are much lighter than protons, their contribution can be ignored.) In reality, however, the ratio is 4:1. Rutherford and others postulated that there must be another type of subatomic particle in the atomic nucleus; the proof was provided by another English physicist, James Chadwick, in 1932. When Chadwick bombarded a thin sheet of beryllium with particles, a very high energy radiation similar to rays was emitted by the metal. Later experiments showed that the rays actually consisted of electrically neutral particles having a mass slightly greater than that of protons. Chadwick named these particles neutrons. The mystery of the mass ratio could now be explained. In the helium nucleus there are two protons and two neutrons, but in the hydrogen nucleus there is only one proton and no neutrons; therefore, the ratio is 4:1. Figure 2.8 shows the location of the elementary particles (protons, neutrons, and electrons) in an atom. There are other subatomic particles, but the electron, the 1 pm 1 1012 m 11013 34 CHAPTER 2 Atoms, Molecules, and Ions If the size of an atom were expanded to that of this sports stadium, the size of the nucleus would be that of a marble. A common non-SI unit for atomic length is the angstrom (Å; 1 Å 100 pm). Proton Neutron Figure 2.8 The protons and neutrons of an atom are packed in an extremely small nucleus. Electrons are shown as “clouds” around the nucleus. cha485 2.3 Atomic Number, Mass Number, and Isotopes 35 proton, and the neutron are the three fundamental components of the atom that are important in chemistry. Table 2.1 shows the masses and charges of these three elementary particles. 2.3 Atomic Number, Mass Number, and Isotopes All atoms can be identified by the number of protons and neutrons they contain. The number of protons in the nucleus of each atom of an element is called the atomic number (Z). In a neutral atom the number of protons is equal to the number of electrons, so the atomic number also indicates the number of electrons present in the atom. The chemical identity of an atom can be determined solely by its atomic number. For example, the atomic number of nitrogen is 7; this means that each neutral nitrogen atom has 7 protons and 7 electrons. Or viewed another way, every atom in the universe that contains 7 protons is correctly named “nitrogen.” The mass number (A) is the total number of neutrons and protons present in the nucleus of an atom of an element. Except for the most common form of hydrogen, which has one proton and no neutrons, all atomic nuclei contain both protons and neutrons. In general, the mass number is given by The number of neutrons in an atom is equal to the difference between the mass number and the atomic number, or (A Z). For example, if the mass number of a particular boron atom is 12 and the atomic number is 5 (indicating 5 protons in the nucleus), then the number of neutrons is 12 5 7. Note that all three quantities (atomic number, number of neutrons, and mass number) must be positive integers, or whole numbers. In most cases atoms of a given element do not all have the same mass. Atoms that have the same atomic number but different mass numbers are called isotopes. For example, there are three isotopes of hydrogen. One, simply known as hydrogen, has one proton and no neutrons. The deuterium isotope has one proton and one neutron, and tritium has one proton and two neutrons. The accepted way to denote the atomic number and mass number of an atom of element X is as follows: mass number A ZX atomic number atomic number number of neutrons mass number number of protons number of neutrons Charge Particle Mass (g) Coulomb Charge Unit Electron* 9.10938 1028 1.6022 1019 1 Proton 1.67262 1024 1.6022 1019 1 Neutron 1.67493 1024 0 0 *More refined experiments have given us a more accurate value of an electron’s mass than Millikan’s. TABLE 2.1 Mass and Charge of Subatomic Particles cha48518_ch0 Thus, for the isotopes of hydrogen, we write As another example, consider two common isotopes of uranium with mass numbers of 235 and 238, respectively: The first isotope is used in nuclear reactors and atomic bombs, whereas the second isotope lacks the properties necessary for these applications. With the exception of hydrogen, isotopes of elements are identified by their mass numbers. Thus these two isotopes are called uranium-235 (pronounced “uranium two thirty-five”) and uranium238 (pronounced “uranium two thirty-eight”). The chemical properties of an element are determined primarily by the protons and electrons in its atoms; neutrons do not take part in chemical changes under normal conditions. Therefore, isotopes of the same element have similar chemistries, forming the same types of compounds and displaying similar reactivities. 235 92U 238 92U 1 1H hydrogen 2 1H deuterium 3 1H tritium 1 1 H 1 2 H 1 3 H.