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Unit 3 I. Structure of the Atom I. Structure of the Atom 1. Many philosophers of ancient times concluded that matter was composed of things such as earth, water, air, and fire. 2. It was also commonly accepted that matter could be endlessly divided into smaller and smaller pieces. 3. Democritus, a Greek philosopher (460-370 B.C.), was the first to propose the idea that matter was not infinitely divisible; he thought matter was made up of tiny individual particles called atomos. I. Structure of the Atom I. Structure of the Atom 4. Aristotle (384-322 B.C.), a Greek philosopher, rejected the atomic “theory” because it did not agree with his own ideas on nature. • The influence of Aristotle was so great and the development of science so primitive that his denial of the existence of atoms went largely unchallenged for 2000 years. I. Structure of the Atom 7. Dalton’s Atomic Theory a. All matter composed of small particles called atoms. b. All atoms of a given element are identical, having the same size, mass, and chemical properties. Atoms of a specific element are different from those of any other element. c. Atoms can’t be created, divided, or destroyed. d. Different atoms combine in simple wholenumber ratios to form compounds. e. In a chemical reaction, atoms are separated, combined, or rearranged. 5. The concept of the atom was revived in the 18th century, but it took another 100 years for significant progress to be made. 6. John Dalton (1766-1844), a teacher in England, revised Democritus’s ideas based on his own scientific research and marked the beginning of the modern atomic theory. I. Structure of the Atom 8. Errors in Dalton’s Atomic Theory a. Atoms ARE divisible into several subatomic particles. b. All atoms of a given element don’t always have identical properties (atoms of an element may have slightly different masses). 9. Atom: the smallest particle of an element that retains the properties of the element 1 I. Structure of the Atom 10. Thomson Model (1904): English physicist J.J. Thomson inferred from his experiments that atoms contained small, negatively charged particles. He thought these “electrons” were evenly embedded throughout a positively charged sphere, much like chocolate chips in a ball of cookie dough. 11. Electron: a negatively charged, fast-moving particle with an extremely small mass that is found in all forms of matter and moves through the empty space surrounding an atom’s nucleus. I. Structure of the Atom 14. Bohr Model (1913): Danish physicist Niels Bohr hypothesized that electrons traveled in fixed orbits around the atom’s nucleus, which contained positively charged particles. James Chadwick, a student of Rutherford, concluded that the nucleus contained positive protons and neutral neutrons. 15. Proton: a subatomic particle in an atom’s nucleus that has a positive charge of 1+. 16. Neutron: a neutral subatomic particle in an atom’s nucleus that has a mass nearly equal to that of a proton. I. Structure of the Atom 18. Quarks: particles of matter that make up protons and neutrons. 19. Atomic Number: the number of protons in an atom 20. Atomic # = # of protons = # of electrons (in a neutral atom) 21. # of neutrons = Mass # – Atomic # 22. Isotopes: atoms with the same number of protons but different numbers of neutrons I. Structure of the Atom 12. Rutherford Model (1911): Another British physicist, Ernest Rutherford, proposed that almost all the mass of an atom – and all its positive charges – were concentrated in a central atomic nucleus surrounded by electrons. 13. Nucleus: the extremely small, positively charged, dense center of an atom that contains positively charged protons, neutral neutrons, and is surrounded by empty space through which one or more negatively charged electrons move. I. Structure of the Atom 17. Electron Cloud Model (1926 – present day): According to the currently accepted model of atomic structure, electrons do not follow fixed orbits but tend to occur more frequently in certain areas around the nucleus at any given time. I. Structure of the Atom 23. Mass Number: used to identify each of the various isotopes of an element; the sum of the number of protons and neutrons in the nucleus; always a whole number (example: carbon-13) 24. The masses of both protons and neutrons are approximately 1.67 x 10-24 g. 2 I. Structure of the Atom 25. Because these numbers are difficult to work with, a different system is used. 26. Scientists assigned the carbon-12 atom a mass of exactly 12 atomic mass units. 27. One atomic mass unit (amu) is defined as 1/12 the mass of a carbon12 atom. I. Structure of the Atom 29. Isotope Notation Mass # Atomic # 107 Ag 47 1- Charge II. Bohr Diagrams 1. A row of elements on the periodic table is called a period. 2. The period # indicates the number of energy levels in an atom of an element in that period. 3. A column of elements on the periodic table is called a group or family. 4. The group # indicates the # of valence electrons (electrons in the outer energy level). I. Structure of the Atom 28. Atomic Mass: the weighted average mass of the isotopes of that element Example: Chlorine exists naturally as a mixture of about 75% chlorine-35 and 25% chlorine-37. Because atomic mass is a weighted average, the chlorine-35 atoms, which exist in greater abundance, have a greater effect in determining the atomic mass. I. Structure of the Atom 30. The presence of a charge on an atom means it has more or fewer electrons than protons. 31. A negative charge means it has more electrons than protons. 32. A positive charge means it has fewer electrons than protons. 33. Ion: An atom with a charge II. Bohr Diagrams 5. To draw a Bohr diagram, write the number of protons (p+) and neutrons (n0) in the nucleus. Then draw circles around the nucleus – one for each energy level. Determine the total number of electrons in the atom (same as the atomic number, unless it’s an ion). Fill in the electrons on the circles, with a maximum of 2 electrons on the first circle, 8 electrons on the second, 18 electrons on the third, and 32 electrons on the fourth. 3 II. Bohr Diagrams 6. Before placing more than 8 electrons on any energy level above the second energy level, you must first place the correct number of valence electrons in the outer energy level. Then go back and fill the lower energy levels with any remaining electrons. III. Changes in the Nucleus 6. Atoms with mass numbers around 60 are the most stable. In order to increase their stability, heavy atoms sometimes fragment into smaller atoms, and light atoms sometimes combine into larger atoms. 7. Nuclear fission: the splitting of a nucleus into fragments 8. Nuclear fusion: the combining of atomic nuclei III. Changes in the Nucleus 1. 2. 3. 4. 5. Strong Force: force that acts between protons and neutrons in an atomic nucleus and keeps them together Nuclear Reaction: a reaction that involves a change in the nucleus of an atom Radioactive Decay: a spontaneous process in which unstable nuclei lose energy by emitting radiation Radioisotopes: isotopes of atoms with unstable nuclei Why do these atoms emit radiation? TO BECOME MORE STABLE III. Changes in the Nucleus 9. Transmutation: the conversion of an atom of one element to an atom of another element 10. Half-life: the time required for one-half of a radioisotope’s nuclei to decay into its products Amount remaining = (Initial Amount)(1/2)n n = # of half-lives III. Changes in the Nucleus 11. Sample Problem: The half-life of iron-59 is 44.5 days. How much of a 2.000 mg sample will remain after 133.5 days? III. Changes in the Nucleus 12. Three types of radiation: a. Alpha: made up of alpha particles, which contain 2 protons and 2 neutrons and have a 2+ charge; equivalent to a helium-4 nucleus b. Beta: made up of fast-moving electrons called beta particles, which have a 1- charge c. Gamma: consists of gamma rays, which are high-energy radiation that possess no mass or charge; usually accompany alpha and beta radiation and account for most of the energy lost during the radioactive process 4 III. Changes in the Nucleus 13. Characteristics of Radiation Radiation Type Symbol Alpha 4 He 2 Beta or Charge 4 2+ 1 1840 0 1- α 0 β -1 Gamma Mass (amu) 0 γ 0 III. Changes in the Nucleus 14. Nuclear Equation: a type of equation that shows the atomic number and mass number of the particles involved 15. Alpha radiation – The alpha decay of radioactive radium-226 into radon-222 is shown below: 226 0 III. Changes in the Nucleus 16. Beta radiation – The beta decay of radioactive carbon-14 into nitrogen-14 is shown below: 14 6 C 14 7 N + 0 -1 18. Positron Emission: a radioactive decay process that involves the emission of a positron from a nucleus; a proton in the nucleus is converted into a neutron and a positron, and the positron is emitted 19. Positron: a particle with the same mass as an electron but opposite in charge 0 1 + 4 2 He 17. Gamma radiation – For example, gamma rays accompany the alpha decay of uranium-238 as shown below: 234 U 92 III. Changes in the Nucleus Rn 86 III. Changes in the Nucleus 238 β 222 Ra 88 Th 90 + 4 2 He + 2 0 0 γ III. Changes in the Nucleus 20. Example of Positron Emission: 26 14 Si 26 Al 13 + 0 1 β β 5 III. Changes in the Nucleus 21. Electron Capture: occurs when the nucleus of an atom draws in a surrounding electron; this electron combines with a proton to form a neutron 22. Example of an Electron Capture: 0 -1 e + 81 Rb 37 81 36 Kr + X-ray photon III. Changes in the Nucleus EXERCISE #1 Alpha decay of iridium-174 EXERCISE #2 Beta decay of platinum-199 EXERCISE #3 Nickel-60, a beta particle, and a gamma ray are the products of the decay of which isotope? III. Changes in the Nucleus EXERCISE #4 Positron emission from sulfur-31 EXERCISE #5 Beryllium-7 undergoes electron capture. EXERCISE #6 The half-life of radium-224 is 3.66 days. What was the original mass of this radioactive isotope if 1.2 mg remains after 10.98 days? 6