Unit Two: Atomic Structure
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1 Unit Two: Atomic Structure TEKS 5: The student understands the historical development of the Periodic Table and can apply its predictive power. (b) use the Periodic Table to identify and explain the properties of chemical families, including alkali metals, alkaline earth metals, halogens, noble gases, and transition metals. TEKS 6: The student knows and understands the historical development of atomic theory. (a) understand the experimental design and conclusions used in the development of modern atomic theory, including Dalton s Postulates, Thomson s discovery of electron properties, Rutherford s nuclear atom, and Bohr s nuclear atom; (b) use isotopic composition to calculate average atomic mass of an element. TEKS 12: The student understands the basic processes of nuclear chemistry. (a) describe the characteristics of alpha, beta, and gamma radiation; Objectives for Subunit 3: Structure of the Atom (Chapter 4) 1) Describe the organization of elements on the periodic table. Identify the period and group on the periodic table. 2) Identify element as metal, metalloid, or non-metal; more specifically an alkali metal, alkaline earth metal, halogen, or noble gas. 3) Identify some of the properties of metals, metalloids and non-metals. 4) Memorize the symbols and names of common elements. This will include the first 36 elements and some other common elements. (Rb, Sr, Cs, Ba, Ag Cd, Sn, Sb, I, Au, Hg, Pb, Bi, U) 5) Know the seven diatomic elements (H 2, O 2, N 2, F 2, Cl 2, Br 2, and I 2 ). 6) Predict the natural physical state of the elements at 25 C. (ie. Silver is solid, nitrogen is gaseous, bromine is liquid.) 7) Summarize Dalton s atomic theory. 8) Distinguish between protons, electrons, and neutrons in terms of their relative masses, charges, and position in atom. 9) Use the atomic number and mass number of an element to find the number of protons, electrons, and neutrons. 10) Use the periodic table to infer the number of valence electrons in an atom and predict the charges of the representative elements. 11) State the octet rule. 12) Describe the formation of cations from metals and anions from nonmetals. 13) Define an atomic mass unit and what is the standard that the relative mass is set. 14) Use the concept of isotopes to explain why the atomic masses of elements are not whole numbers. 15) Calculate the average atomic mass of an element from isotope data. 16) Describe the contributions and experiments that Thomson, Millikan, and Rutherford made to the development of atomic theory. 17) Compare and contrast the models of the atom proposed by Dalton, Thomson and Rutherford. 18) Explain the relationship between unstable nuclei and radioactive decay. 19) Explain why certain nuclei are unstable according to the neutron-to-proton ratio and relate the band of stability to this ratio. 20) Characterize alpha, beta, and gamma radiation in terms of mass, charge and penetrating power. Describe the experiment used to determine the particles. Vocabulary Periodic table noble gas atomic number radioactivity period diatomic element isotope radiation group conductivity mass number nuclear reaction metal ductility atomic mass unit (amu) radioactive decay metalloid luster atomic mass alpha radiation semiconductor Dalton s Atomic Theory ion alpha particle nonmetal atom cation nuclear equation alkali metal cathode ray anion beta radiation alkaline earth metal electron nucleus beta particle transition metal proton natural abundance gamma ray halogen neutron relative abundance malleability Unit Two Theme: Atoms are the fundamental building blocks of matter.
2 Unit Two Essential Questions: I. What are the similarities and differences of the atomic models of Democritus, Aristotle, and Dalton? II. How was Dalton s theory used to explain the conservation of mass? III. What is an atom? IV. How con the subatomic particles be distinguished in terms of relative charge and mass? V. Where are the locations of the subatomic particles within the structure of the atom? VI. How is the atomic number used to determine the identity of an atom? VII. Why are atomic masses not whole numbers? VIII. Given the mass number and atomic number, how are the number of electrons, protons, and neutrons in an atom calculated? IX. How does changing the subatomic particles effect the atom? X. What is the relationship between unstable nuclei and radioactive decay? XI. How are alpha, beta, and gamma radiation characterized in terms of mass and charge? Unit Two Basic Questions: 1. What is an element? How many elements are presently known? How many of these occur naturally, and how many are human-made? Which elements are most abundant on the earth? 2. State the ideas that Democritus had about matter? Why was it so difficult for him to defend his ideas? 3. State Dalton s Atomic Theory. There are six ideas involved. 4. What part of the atom did J.J. Thompson discover and what was the experiment that provided the evidence? Provide a sketch. 5. What model of the atom did Rutherford propose and what was the experiment that provided evidence for this model? Provide a sketch. 6. What are the three fundamental particles that compose all atoms? Indicate the relative electrical charge and relative mass of each of these particles. Where is each type of particle found in the atom? 7. What part of the atom identifies the element? 8. Explain why atoms are electronically neutral. 9. What particles account for most of the atom s mass? Explain. 10. What is an atomic mass unit (amu) and what is the standard by which atomic mass is determined? 11. What are isotopes? Does changing the number of neutrons affect the way the atom will behave chemically. Explain. 12. How can you determine the number of neutrons in an atom if its mass number and its atomic number are known? 13. What does the subscript and the superscript represent in an isotope notation, also known as nuclear notation? 14. Does the existence of isotopes contradict part of Dalton s original atomic theory? Explain. 15. What is radioactive decay? 16. Why are some atoms radioactive? 17. Define the three types of radioactive decay particles and provide a sketch of the experiment the provided evidence for the particles. 18. Are most elements found in nature in the elemental or in the combined form? Why? Name several elements that are usually found in the elemental form. 19. Give the names of some of the families of elements in the periodic table. 20. Which general area of the periodic table contains the metallic elements? Which general area contains the nonmetallic elements? 21. What are ions? How are ions formed from atoms? To what do the terms cation and anion refer? 22. Write the symbol and atomic number for each of the following elements: magnesium, tin, lead, sodium, hydrogen, chlorine, and silver. 23. Write the name and atomic number for each of the following elements: a. He b. B c. Se d. Ba e. P f. Sr 24. Write the name and chemical symbol for each of the following elements. a. 19 b. 12 c. 1 d. 6 e. 82 f Indicate the number of protons, neutrons, and electrons in isolated atoms having the following nuclear symbols: a. 4 He b. 37 Cl c. 40 Ca 26. What simple ion does each of the following elements most commonly form? a. Mg b. F c. Ba d. Na e. O f. Cl 27. For each of the following simple ions, indicate the number of protons and electrons the ion contains. a. K 1+ b. Br 1- c. Na 1+
3 Unit Two Standard Questions: 28. Why do the symbols for some elements seem to bear no relationship to the name for the element? Give several examples and explain. 29. Explain how Dalton s theory of the atom and the conservation of mass are related. 30. What parts of Dalton s Theory of the atom were later determined to not be completely correct? Explain. 31. Discuss the points of Thomson s cathode-ray experiment that provide evidence to support the discovery of the electron. Why is Thomson s model of the atom called the plum pudding model? 32. Discuss the points of Rutherford s model for the nuclear atom and how he tested this model. Based on his experiments, how did Rutherford envision the structure of the atom? 33. How did Rutherford s model for atomic structure differ from Thomson s plum pudding model? 34. Which of the subatomic particles is responsible for the chemical behavior of a given type of atom? Why? 35. To what do the atomic number and the mass number of an isotope refer? How are specific isotopes indicated symbolically? Give an example and explain. How are isotopes named to distinguish one isotope of the same element from another? Give an example. 36. Explain how to calculate atomic mass. What is the difference between atomic mass and mass number? 37. What is the difference between percent abundance and relative abundance? 38. In terms of subatomic particles, how is a cation related to the atom form which it is formed? An anion? Does the nucleus of an atom change when an atom is converted into an ion? 39. Explain how unstable atoms gain stability. 40. What quantities are conserved when balancing a nuclear reaction? 41. Compare and contrast nuclear and chemical reactions. 42. What are valence electrons and how are valence electrons determined from the periodic table? 43. What is the octet rule and how is it related to the formation of ions? Unit Two Challenge Questions: 44. Compare and contrast Dalton s ideas of matter to those of Democritus. Why did Dalton s ideas last? 45. Describe how the Millikan oil-drop experiment was used to determine the charge of an electron. Sketch a picture of the experiment. What is the actual charge on a single electron? What is the relative charge of an electron? 46. An element has three naturally occurring isotopes. What other information must you know in order to calculate the elements atomic mass? 47. Complete the following table: Isotope Atomic Mass Number of Number of Number of Element Notation Number Number Protons Neutrons Electrons Name Sn Determine the atomic mass of Titanium from the following information: Table of Relative Abundances of Titanium Isotopes Element Mass (amu) Abundance (%) Titanium Titanium Titanium Titanium Titanium Boron-10 and boron-11 are the naturally occurring isotopes of elemental boron. If boron has an atomic mass of amu, which isotope occurs in greater abundance? Explain your reasoning. 50. Explain how the existence of isotopes is related to the fact that atomic masses are not whole numbers.
4 Structure of the Atom 30. Summarize Dalton's atomic theory. 31. Six atoms of Element A combine with eight atoms of Element B to produce six compound particles. How many atoms of Elements A and B does each particle contain? Are all of the atoms used to form compounds? 32. Compare and contrast Thomson's plum pudding atomic model with Rutherford's nuclear atomic model. 33. Compare the relative charge, mass and location of each of the subatomic particles. 34. Which subatomic particle identifies an atom as that of a particular element? 35. How do isotopes of a given element differ? How are they similar? 36. How is an atom's atomic number related to its number of protons? To its number of electrons? 37. How is the mass number related to the number of protons and neutrons an atom has? 38. How can you determine the number of neutrons in an atom if its mass number and its atomic number are known? 39. How is the charge of an atom related to the electrons and protons of an atom? How is a negative ion determined and what is it called? How is a positive ion determined and what is it called? 40. Fill the following table: Isotope Atomic Mass Number of Number of Number of Element Notation Number Number Protons Neutrons Electrons Name Sn How many electrons, protons, and neutrons are contained in each atom? a) gallium-69 b) fluorine-23 c) titanium-48 d) tantalum Determine the atomic mass of sulfur given the following naturally occurring isotopes: 32 S ( amu, 94.99%); 33 S ( amu, 0.75%); 34 S ( amu, 4.25%); and 36 S ( amu, 0.01%). 43. Use the table to answer the following questions: Table of Relative Abundances of Zirconium Isotopes Element Mass (amu) Abundance (%) Zirconium Zirconium Zirconium Zirconium Zirconium a) What is the mass number of each zirconium isotope? b) Compute the number of protons and neutrons for each zirconium isotope. c) Does the number of protons or neutrons remain the same for all isotopes? Explain. d) Base on the relative abundances of each isotope, predict to which isotope's mass the average atomic mass of zirconium is going to be closest. e) Calculate the weighted average atomic mass of zirconium. Nuclear Chemistry 44. Classify each of the following as a chemical reaction, a nuclear reaction, or neither. a) Thorium emits a beta particle. b) Two atoms share electrons to form a bond. c) A sample of pure sulfur emits heat energy as it slowly cools. d) A piece of iron rusts. 45. What is radioactive decay? 46. Why are some atoms radioactive? What is the primary factor that determines whether a nucleus is stable or unstable? 47. Write the symbols used to denote alpha, beta, and gamma radiation and give their mass and charge. 48. Write the completed and balanced nuclear equation: [Name the daughter nuclides] a) U 4 2He + b) 9 4Be C + 1 0n 231 c) 91 Pa Np d) 35 17Cl S + 4 2He e) Plutonium-256 goes through three beta decays and two alpha decays. 49. Polonium-208 goes through natural alpha decay, a beta decay and another alpha decay. This isotope has a half-life of 59.3 seconds.
5 a) Write the correct nuclear equation for this isotope identifying the resulting daughter nuclide. b) If you have a 12.0 gram sample of Po-208, how long will it be until you have a 0.75 g sample of the nuclide. 50. Write balanced equations for each of the processes described below. a) Chromium-51, which targets the spleen and is used as a tracer in studies of red blood cells, decays by electron capture. b) Iodine-131, used to treat hyperactive thyroid glands, decays by producing a β particle. c) Phosphorus-32, which accumulates in the liver, decays by β-emission. d) Uranium-235, which is used in atomic bombs, decays initially by α-emission. 51. Write an equation describing the radioactive decay of each of the following nuclides. (The particle produced is shown in parentheses, except for electron capture, where an electron is a reactant.) a) 68 Ga (e - capture) b) 62 Cu (positron) c) 212 Fr (α) d) 129 Sb (β) 52. In each of the following nuclear reactions, supply the missing particle. a) 73 Ga 73 Ge + b) 192 Pt 188 Os + c) 205 Bi 205 Pb + d) 241 Cm Am 53. One type of commercial smoke detector contains a minute amount of radioactive americium-241, which decays by α- emission. The α particles ionize molecules in the air, allowing it to conduct an electric current. When smoke particles enter, the conductivity of the air is changed and the alarm buzzes. a) Write the equation for the decay of 241 Am by alpha-emission. b) The complete decay of 241 Am involves successively α, α, β, α, α, β, α, α, α, β, α, and β emission. What is the final stable nucleus produced in this decay series? 54. In 1994 it was proposed that element 106 be named seaborgium, Sg, in honor of Glenn T. Seagorg, discoverer of the transuranium elements. a) 263 Sg was produced by the bombardment of 249 Cf with a beam of 18 O nuclei. Complete and balance an equation for this nuclear transmutation. b) 263 Sg decays by α-emission. What is the other product resulting from this decay? 55. Many elements have been synthesized by bombarding relatively heavy atoms with high-energy particles in particle accelerators. Complete the following transmutation nuclear reactions, which have been used to synthesize elements. 4 a) _ He 97Bk 0n b) 92U 6C _ 60 n c) 98Cf _ 105Db 40n d) 98 Cf 5B 103Lr _ 56. Technetium-104 has a half-life of 18.0 min. How much of a g sample remains after 90.0 minutes have passed? 57. What is the half-life of radon-222 if a sample initially contains 150 mg and only 18.7 mg after 11.4 days? 58. Compare and contrast nuclear fission and nuclear fusion reaction. 59. Explain how nuclear fission can be used to generate electric power. 60. Explain the purpose of control rods in a nuclear reactor. Explain the purpose of the moderator in a nuclear reactor. 61. Describe what is meant by the terms critical mass and subcritical mass.
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