The Nuclear Atom

Atoms are composed of three sub-atomic particks protons, neutrons and ekctrons. 

Paradigm shifts - the sub-atomic particle theory of matter 

Atomism, the concept that matter is composed of atoms, is thousands of years old. The idea was founded in philosophical reasoning rather than experimentation. The earliest references to atomism date back to India in the sixth century bce. The earliest references in the Western World emerged a century later in Ancient Greece, from Leucippus and his student Democritus. 

In 1803, the British chemist John Dalton (1766-1844) developed an atomic theory to explain why chemical elements react in simple proportions by mass. He proposed that each chemical element consisted of identical atoms and that these atoms could bond together to form chemical compounds. It is not clear whether Dalton was aware of previous ancient ideas about atoms, but his interest in atoms was strongly influenced by the experiments he performed on gases. He did not prove the existence of atoms he simply demonstrated that his atomic theory was consistent with experimental data. 

Sub-atomic particle Symbol Relative mass Relative charge Nuclide notation 

How are the subatomic particles arranged to form an atom This puzzled scientists for many years until a British physicist, Ernest Rutherford, and his colleagues discovered the answer in 1911. Their experiments led them to conclude that nearly all the mass of an atom is packed into a tiny, positive-charged body called the nucleus, centered in a very much larger cloud of electrons. The atom was seen to be mostly empty space  

Protons and neutrons, the two massive subatomic particles, make up the nucleus and are responsible for nearly all the mass of the atom. The size of the enveloping cloud of electrons is from 10,000 to 20,000 times the size of the nucleus, and the size of this electron cloud represents the size of the atom. Although the electron cloud of an iron atom, Fe, is immensely larger than the nucleus, the mass of the nucleus is about 4,000 times greater than the mass of all the electrons surrounding it. 

7Democritus, Greek philosopher, bom Abdera, Thrace (the eastern Balkans) ca. 470 bc. Died ca. 

It was perhaps Thomson who first suggested a specific structure for the atom in terms of subatomic particles. His plum pudding model (ca. 1900), which placed electrons in a sea of positive charge, like raisins in a pudding., accorded with the then-known facts in evidently permitting electrons to be removed under the influence of an electric potential. The modem picture of the atom as a positive nucleus with extranuclear electrons was proposed by Rutherford13 in 1911. It arose from 

10Jean Perrin, bom Lille, France, 1870. Ph.D. Ecole Normal Superieure, Paris. Professor University of Paris. Nobel Prize in physics 1923. Died New York, 1942. 

1 Wilhelm Friedrich Ostwald, German chemist, bom Riga, Latvia, 1853. Ph.D. Dorpat, Estonia. Professor Riga, Leipzig. A founder of physical chemistry, opponent of the atomic theory till convinced by the work of Einstein and Perrin. Nobel Prize in chemistry 1909. Died near Leipzig, 1932. 

12Sir Joseph John Thomson, bom near Manchester, 1856. Professor, Cambridge. Nobel Prize in physics 1906. Knighted 1908. Died Cambridge, 1940. 

This number of levels is really doubled, however, by the electron spin. An electron has an intrinsic permanent magnetism, and associated with it a permanent angular momentum of magnitude %h/2ir. This can bo oriented in either of two opposite directions, giving a component of ih/2tt along a fixed direction. This, as will be seen, is in harmony with the space quantization just described, for the special case Z = For each stationary state of an electron neglecting spin, we can have the two possible orientations of the spin, so that actually an s level has two 

Z0 is called a shielding constant. It measures the effect of the other electrons in reducing the nuclear attraction for the electron in question. It is a function of n and Z, increasing from practically zero for the lowest n values to a value only slightly less than Z for the outermost orbits within the atom. For levels outside the atom, on the other hand, the energy is given approximately by 

Fig XXI-1.— Energies of electrons in the copper atom, in Rydberg units, us a function of principal quantum number n. Eneigies are shown on a logai ithmie scale. The energies in the hydrogen atom are shown for comparison. 

It has been mentioned that the region occupied by the electron s orbit increases in volume, as the binding energy becomes less or as the quantum number n increases. For our later use in studying the sizes of atoms, it is useful to know the size of the orbit quantitatively. These sizes are not definitely determined, for the electron is sometimes found at one point, sometimes at another, in a given stationary state, and all we can give is the distance from the nucleus at which there is the greatest 

We may expect that the radius of the atom, if that expression has a meaning, will be of the order of magnitude of the radius of the largest orbit ordinarily occupied by an electron in the neutral atom. In the case of copper this is the 4s orbit, while in the copper ion it is the 3d. In the next section we tabulate such quantities for the atoms, and in later chapters we shall find these radii of interest in connection with the dimensions of atoms as determined in other ways. 

Electric charge is more fully defined in Secticn 4.4. For now, think of it as an inherent property of eleotrons that causes them to interact with other charged particles. 

By the end of the nineteenth century, scientists were convinced that matter was composed of atoms, the permanent, indestructible building blocks from which all substances are constructed. However, an English physicist named J. J. Thomson (1856-1940) complicated the picture by discovering an even smaller and more fundamental particle called the electron. Thomson discovered that electrons are negatively charged, that they are much smaller and lighter than atoms, and that they are uniformly present in many different kinds of substances. The indestructible building block called the atom could apparently be chipped.  

Most of the atom s mass and all of its positive charge are contained in a small core called the nucleus. 

Most of the volume of the atom is empty space through which the tiny, negatively charged electrons are dispersed. 

There are as many negatively charged electrons outside the nucleus as there are positively charged particles (protons) inside the nucleus, so that the atom is electrically neutral. 

The large-angle scattering greatly puzzled Rutherford. As he commented some years later, this observation was about as credible as if you had fired a 15-inch shell at a piece of tissue paper and it came back and hit you. By 1911, though, Rutherford had an explanation. He based his explanation on a model of the atom known as the nuclear atom and having these features  

Most of the mass and all of the positive charge of an atom are centered in a very small region called the nucleus. The remainder of the atom is mostly empty space. 

The magnitude of the positive charge is different for different atoms and is approximately one-half the atomic weight of the element. 

The telescope travels in a circular track around an evacuated chamber containing the metal foil. Most a particles pass through the metal foil undeflected, but some are deflected through large angles. 

TABLE 5.2 Electrical Charge and Relative Mass of Electrons, Protons, and Neutrons 

Particle Symbol Relative electrical charge Actual mass (g) 

The mass of a helium atom is 6.65 X 10 g. How many atoms are in a 4.0-g sample of helium  

The discovery that positively charged particles are present in atoms came soon after the discovery of radioactivity by Henri Becquerel (1852-1908) in 1896. Radioactive elements spontaneously emit alpha particles, beta particles, and gamma rays from their nuclei (see Chapter 18). 

When we speak of the mass of an atom, we are referring primarily to the mass of the nucleus. The nucleus contains all the protons and neutrons, which represent more than 99.9% of the total mass of any atom (see Table 5.1). By way of illustration, the largest number of electrons known to exist in an atom is 116. The mass of even 116 electrons is only about 1/17 of the mass of a single proton or neutron. The mass of an atom therefore is primarily determined by the combined masses of its protons and neutrons. 

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Scientists have developed a highly sophisticated view of the structure of the atom. The currently accepted model is called the nuclear atom. We shall present it without trying to show immediately all of the experimental evidence that led to this particular model. Rest assured, though, that every feature of the nuclear atom picture rests upon experimental evidence, as we shall see in Chapter 14. 

The most stable state of the atom would be expected to be the one in which the atom has the lowest energy. Bohr reasoned that since we observe that the nuclear atom does exist then it must be a fundamental fact of nature that an atom can exist in its most stable state indefinitely. Even though this fact could not be rationalized (remember, the earlier laws of physics predicted the atom should collapse) it had to be accepted because it was a result of experiments. 

Read about the gold foil experiment in your textbook. Describe the plum-pudding atomic model. How did the gold foil experiment show the plumpudding model to be in error Describe the nuclear atomic model that replaced the plumpudding model. 

The first application of quantum theory to a problem in chemistry was to account for the emission spectrum of hydrogen and at the same time explain the stability of the nuclear atom, which seemed to require accelerated electrons in orbital motion. This planetary model is rendered unstable by continuous radiation of energy. The Bohr postulate that electronic angular momentum should be quantized in order to stabilize unique orbits solved both problems in principle. The Bohr condition requires that... 

Rutherford s attitude toward chemistry was stereotyped by his jokes and barbs occasionally directed at his chemical colleagues. The later Manchester physicist P. M. S. Blackett recounted the famous crack, "All science is either physics or stamp collecting, "63 and it was said that Rutherford chafed at receiving the 1908 Nobel Prize in chemistry, rather than in Physics. In a lecture in which he described his theory of the nuclear atom, he joked that the "nucleus is a round, hard objectjust like Professor Perkins head."64 However, Rutherford expressed great respect for his chemist collaborator Frederick Soddy and for other chemists, as well. 

The concept of the atom as the smallest particle of matter (from the Greek word for indivisible) was promulgated by John Dalton about 1803. Within about a century and a quarter of scientific investigation which will be briefly described in this chapter, this concept yielded the idea of the periodic table and the understanding of the periodic table including the nuclear atom, the concept of isotopes, and the discovery of the majority of the isotopes which are used in the studies of the isotope effects. It is appropriate to point out that this book deals with the study of the effect of isotopic substitution on the physical and chemical properties of molecular (or atomic) systems. The book does not deal with the use of isotopes as tracers, a use which usually depends on the assumption that isotope effects are small and can be ignored in tracer studies. 

It was quite the most incredible event . Quoted in G. K. T. Conn and H. D. Turner, The Evolution of the Nuclear Atom (London Iliffe Books, 1965), 136. 

A series of episodes in the historical development of our view of chemical atoms are presented. Emphasis is placed on the key observations that drove chemists and physicists to conclude that atoms were real objects and to envision their stracture and properties. The kinetic theory of gases and measmements of gas transport yielded good estimates for atomic size. The discovery of the electrorr, proton and neutron strongly irtfluenced discttssion of the constitution of atoms. The observation of a massive, dertse nucleus by alpha particle scattering and the measrrrement of the nuclear charge resrrlted in an enduring model of the nuclear atom. The role of optical spectroscopy in the development of a theory of electronic stracture is presented. The actors in this story were often well rewarded for their efforts to see the atoms. 

Molybdo-vanadoarsenates.—A number of compounds have been described which are analogous to the molybdo-vanadophosphates described above, and which contain arsenic for the nuclear atom of the complex anion. In many cases these compounds approximate to the... 

B.2 The Nuclear Atom B.3 Neutrons B.4 Isotopes B.5 The Organization of the Elements... 

In the nuclear atom, all the positive charge and almost all the mass is concentrated in the tiny nucleus, and the negatively charged electrons surround the nucleus. The atomic number is the number of protons in the nucleus there is an equal number of electrons outside the nucleus. 

As soon as we start this journey, we encounter an extraordinary feature of our world. When Rutherford proposed the nuclear atom (Section B), he expected to be able to use classical mechanics, the laws of motion proposed by Newton in the seventeenth century, to describe its electronic structure. After all, classical mechanics had been tremendously successful for describing visible objects such as balls and planets. However, it soon became clear that classical mechanics fails when applied to electrons in atoms. New laws, which came to be known as quantum mechanics, were developed in the early part of the twentieth century. 

Neils Bohr (1885-1962) proposed an orbital model of the nuclear atom in which electrons in an atom moved around the nucleus, just as planets move around the sun. 

Rutherford published the results of these scattering experiments in mid-1909, and it seemed as if publication of the discovery of the nuclear atom would soon follow. But the plum pudding model remained the working model of the atom. Through the rest of 1909 and most of 1910, Rutherford pondered. 

Figure 4.2 The nuclear atom of Rutherford. The electrons were distributed within a sphere surrounding the massive nucleus. No electron orbits were specified.

Ernest Rutherford (1871-1937) was born on a farm in New Zealand. In 1895 he placed second in a scholarship competition to attend Cambridge University but was awarded the scholarship when the winner decided to stay home and get married. As a scientist in England, Rutherford did much of the early work on characterizing radioactivity. He named the a and /3 particles and the g ray and coined the term half-life to describe an important attribute of radioactive elements. His experiments on the behavior of a particles striking thin metal foils led him to postulate the nuclear atom. He also invented the name proton for the nucleus of the hydrogen atom. He received the Nobel Prize in chemistry in 1908. 

Soon after the development of the concept of the nuclear atom, some idea was gained as to the way in which a proton and an electron are combined to form a hydrogen atom In 1913 Niels Bohr (1885- ),... 

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