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3.1 Figure 1 In Dalton’s atomic model, an atom is a solid sphere, similar to a billiard ball. This simple model is still used today to represent the arrangement of atoms in molecules. DID YOU KNOW ? William Crookes (1832–1919) William Crookes was the eldest of sixteen children and inherited his father’s fortune, made in real estate. This enabled him to lead a leisurely life, and also to conduct scientific research in many areas of chemistry and physics. Crookes is best known for his cathode ray tube, which was made possible by his improvements to the vacuum pump and Volta’s invention of the electric cell. His vacuum techniques later made mass production of the light bulb practical. Early History of Atomic Theories The history of atomic theories is full of success and failure stories for hundreds of chemists. In textbooks such as this one, only the success of a few is documented. However, the success of these chemists was often facilitated by both the success and failure of many others. Recall that by the use of deductive logic the Greeks (for example, Democritus) in about 300 B.C. hypothesized that matter cut into smaller and smaller pieces would eventually reach what they called the atom — literally meaning indivisible. This idea was reintroduced over two thousand years later by an English chemist/schoolteacher named John Dalton in 1805. He re-created the modern theory of atoms to explain three important scientific laws — the laws of definite composition, multiple proportions, and conservation of mass. The success of Dalton’s theory of the atom was that it could explain all three of these laws and much more. Dalton’s theory was that the smallest piece of matter was an atom that was indivisible, and that an atom was different from one element to another. All atoms of a particular element were thought to be exactly the same. Dalton’s model of the atom was that of a featureless sphere — by analogy, a billiard ball (Figure 1). Dalton’s atomic theory lasted for about a century, although it came under increasing criticism during the latter part of the 1800s. SUMMARY Creating the Dalton Atomic Theory (1805) Table 1 Key experimental work Theoretical explanation Atomic theory Law of definite composition: elements combine in a characteristic mass ratio Each atom has a particular combining capacity. Law of multiple proportions: there may be more than one mass ratio Some atoms have more than one combining capacity. Matter is composed of indestructible, indivisible atoms, which are identical for one element, but different from other elements. Law of conservation of mass: total mass remains Atoms are neither created nor destroyed constant in a chemical reaction. The Thomson Atomic Model The experimental studies of Svante Arrhenius and Michael Faraday with electricity and chemical solutions and of William Crookes with electricity and vacuum tubes suggested that electric charges were components of matter. J. J. Thomson’s quantitative experiments with cathode rays resulted in the discovery of the electron, whose charge was later measured by Robert Millikan. The Thomson model of the atom (1897) was a hypothesis that the atom was composed of electrons (negative particles) embedded in a positively charged sphere (Figure 2(a)). Thomson’s research group at Cambridge University in England used mathematics to predict the uniform three-dimensional distribution of 162 Chapter 3 NEL Section 3.1 the electrons throughout the atom. The Thomson model of the atom is often communicated by using the analogy of a raisin bun, with the raisins depicting the electrons and the bun being the positive material of his atom (Figure 2(b)). (a) SUMMARY INVESTIGATION 3.1.1 The Nature of Cathode Rays (p. 209) The discovery of cathode rays led to a revision of the Dalton atomic model. What are their properties? Figure 2 (a) In Thomson’s atomic model, the atom is a positive sphere with embedded electrons. (b) This model can be compared to a raisin bun, in which the raisins represent the negative electrons and the bun represents the region of positive charge. (b) Creating the Thomson Atomic Theory (1897) Table 2 Key experimental work Theoretical explanation Atomic theory Arrhenius: the electrical nature of chemical solutions Atoms may gain or lose electrons to form ions in solution. Faraday: quantitative work with electricity and solutions Particular atoms and ions gain or lose a specific number of electrons. Crookes: qualitative studies of cathode rays Electricity is composed of negatively charged particles. Matter is composed of atoms that contain electrons (negative particles) embedded in a positive material. The kind of element is characterized by the number of electrons in the atom. Thomson: quantitative studies of cathode rays Electrons are a component of all matter. Millikan: charged oil drop experiment Electrons have a specific fixed electric charge. DID YOU KNOW ? Rutherford Quotes • “You know it is about as incredible as if you fired a 350-mm shell at a piece of tissue paper and it came back and hit you.” • “Now I know what the atom looks like.” 1911 • The electrons occupy most of the space in the atom, “like a few flies in a cathedral.” • The notion that nuclear energy could be controlled is “moonshine.” 1933 The Rutherford Atomic Theory One of Thomson’s students, Ernest Rutherford (Figure 3), eventually showed that some parts of the Thomson atomic theory were not correct. Rutherford developed an expertise with nuclear radiation during the nine years he spent at McGill University in Montreal. He worked with and classified nuclear radiation as alpha (), beta (), and gamma () — helium nuclei, electrons, and high-energy electromagnetic radiation from the nucleus, respectively. Working with his team of graduate students he devised an experiment to test the Thomson model of the atom. They used radium as a source of alpha radiation, which was directed at a thin film of gold. The prediction, based on the Thomson model, was that the alpha particles should be deflected little, if at all. When some of the alpha particles were deflected at large angles and even backwards from the foil, the prediction NEL Figure 3 Rutherford’s work with radioactive materials at McGill helped prepare him for his challenge to Thomson’s atomic theory. Atomic Theories 163 ACTIVITY 3.1.1 Rutherford’s Gold Foil Experiment (p. 210) Rutherford’s famous experiment involved shooting “atomic bullets” at an extremely thin sheet of gold. You can simulate his experiment. Prediction was shown to be false, and the Thomson model judged unacceptable (Figure 4). Rutherford’s nuclear model of the atom was then created to explain the evidence gathered in this scattering experiment. Rutherford’s analysis showed that all of the positive charge in the atom had to be in a very small volume compared to the size of the atom. Only then could he explain the results of the experiment (Figure 5). He also had to hypothesize the existence of a nuclear (attractive) force, to explain how so much positive charge could occupy such a small volume. The nuclear force of attraction had to be much stronger than the electrostatic force repelling the positive charges in the nucleus. Even though these theoretical ideas seemed far-fetched, they explained the experimental evidence. Rutherford’s explanation of the evidence gradually gained widespread acceptance in the scientific community. alpha particles metal foil SUMMARY Evidence alpha particles Creating the Rutherford Atomic Theory (1911) Table 3 metal foil Figure 4 Rutherford’s experimental observations were dramatically different from what he had expected based on the Thomson model. Key experimental work Theoretical explanation Atomic theory Rutherford: A few positive alpha particles are deflected at large angles when fired at a gold foil. The positive charge in the atom must be concentrated in a very small volume of the atom. Most materials are very stable and do not fly apart (break down). A very strong nuclear force holds the positive charges within the nucleus. Rutherford: Most alpha particles pass straight through gold foil. Most of the atom is empty space. An atom is composed of a very tiny nucleus, which contains positive charges and most of the mass of the atom. Very small negative electrons occupy most of the volume of the atom. Protons, Isotopes, and Neutrons nucleus atom Figure 5 To explain his results, Rutherford suggested that an atom consisted mostly of empy space, explaining why most of the alpha particles passed nearly straight through the gold foil. proton ( 01p or p+) a positively charged subatomic particle found in the nucleus of atoms 164 Chapter 3 The Thomson model of the atom (1897) included electrons as particles, but did not describe the positive charge as particles; recall the raisins (electrons) in a bun (positive charge) analogy. The Rutherford model of the atom (1911) included electrons orbiting a positively charged nucleus. There may have been a hypothesis about the nucleus being composed of positively charged particles, but it was not until 1914 that evidence was gathered to support such a hypothesis. Rutherford, Thomson, and associates studied positive rays in a cathode ray tube and found that the smallest positive charge possible was from ionized hydrogen gas. Rutherford reasoned that this was the fundamental particle of positive charge and he named it the proton, meaning first. (Again Rutherford showed his genius by being able to direct the empirical work and then interpret the evidence theoretically.) By bending the hydrogen-gas positive rays in a magnetic field they were able to determine the charge and mass of the hypothetical proton. The proton was shown to have a charge equal to but opposite to that of the electron and a mass 1836 times that of an electron. All of this work was done in gas discharge tubes that evolved into the version of the mass spectrometer (Figure 6) developed by Francis Aston during the period 1919–1925. Evidence from radioactivity and mass spectrometer investigations falsified Dalton’s theory that all atoms of a particular element were identical. The evidence indicated that NEL Section 3.1 magnet slit gas discharge tube sample inlet slit detector magnet beam of positive ions there were, for example, atoms of sodium with different masses. These atoms of different mass were named isotopes, although their existence could not yet be explained. Later, James Chadwick, working with Rutherford, was bombarding elements with alpha particles to calculate the masses of nuclei. When the masses of the nuclei were compared to the sum of the masses of the protons for the elements, they did not agree. An initial hypothesis was that about half of the mass of the nucleus was made up of proton–electron (neutral) pairs. However, in 1932 Chadwick completed some careful experimental work involving radiation effects caused by alpha particle bombardment. He reasoned that the only logical and consistent theory that could explain these results involved the existence of a neutral particle in the nucleus. According to Chadwick, the nucleus would contain positively charged protons and neutral particles, called neutrons. The different radioactive and mass properties of isotopes could now be explained by the different nuclear stability and different masses of the atom caused by different numbers of neutrons in the nuclei of atoms of a particular element. SUMMARY Figure 6 A mass spectrometer is used to determine the masses of ionized particles by measuring the deflection of these particles as they pass through the field of a strong magnet. isotope ( AZ X) a variety of atoms of an element; atoms of this variety have the same number of protons as all atoms of the element, but a different number of neutrons neutron ( 01n or n) a neutral (uncharged) subatomic particle present in the nucleus of atoms Rutherford Model • An atom is made up of an equal number of negatively charged electrons and postively charged protons. • Most of the mass of the atom and all of its positive charge is contained in a tiny core region called the nucleus. • The nucleus contains protons and neutrons that have approximately the same mass. • The number of protons is called the atomic number (Z). • The total number of protons and neutrons is called the mass number (A). NEL Atomic Theories 165 SUMMARY Creating the Concepts of Protons, Isotopes, and Neutrons Table 4 Key experimental work Theoretical explanation Atomic theory Rutherford (1914): The lowest charge on an ionized gas particle is from the hydrogen ion The smallest particle of positive charge is the proton. Soddy (1913): Radioactive decay suggests different atoms of the same element Isotopes of an element have a fixed number of protons but varying stability and mass. Aston (1919): Mass spectrometer work indicates different masses for some atoms of the same element The nucleus contains neutral particles called neutrons. Atoms are composed of protons, neutrons, and electrons. Atoms of the same element have the same number of protons and electrons, but may have a varying number of neutrons (isotopes of the element). Radiation is produced by bombarding elements with alpha particles. Section 3.1 Questions Understanding Concepts 1. Summarize, using labelled diagrams, the evolution of atomic theory from the Dalton to the Rutherford model. 2. Present the experimental evidence that led to the Rutherford model. 3. How did Rutherford infer that the nucleus was (a) very small (compared to the size of the atom)? (b) positively charged? 4. (a) State the experimental evidence that was used in the discovery of the proton. (b) Write a description of a proton. 5. (a) State the experimental evidence that was used in the discovery of the neutron. (b) Describe the nature of the neutron. Applying Inquiry Skills 6. What is meant by a “black box” and why is this an appro- priate analogy for the study of atomic structure? 7. Theories are often created by scientists to explain scientific laws and experimental results. To some people it seems strange to say that theories come after laws. Compare the scientific and common uses of the term “theory.” 8. What is the ultimate authority in scientific work (what kind of knowledge is most trusted)? 166 Chapter 3 Making Connections 9. State some recent examples of stories in the news media that mention or refer to atoms. 10. Describe some contributions Canadian scientists and/or scientists working in Canadian laboratories made to the advancement of knowledge about the nature of matter. GO www.science.nelson.com Extension 11. Rutherford’s idea that atoms are mostly empty space is retained in all subsequent atomic theories. How can solids then be “solid”? In other words, how can your chair support you? Why doesn’t your pencil go right through the atoms that make up your desk? 12. When you look around you, the matter you observe can be said to be made from electrons, protons, and neutrons. Modern scientific theories tell us something a little different about the composition of matter. For example, today protons are not considered to be fundamental particles; i.e., they are now believed to be composed of still smaller particles. According to current nuclear theory, what is the composition of a proton? Which Canadian scientist received a share of the Nobel Prize for his empirical work in verifying this hypothesis of sub-subatomic particles? NEL