spins. As shown in the following table, the sublevels s, p, d, and f have 1, 3, 5, and 7 available orbitals, respectively.

Math Tutor The arrangement of elements in the periodic table reflects the arrangement of electrons in an atom. Each period begins with an atom that has an electron in a new energy level and with the exception of the first period, each period ends with an atom that has a filled set of p orbitals. 1s 2s Writing Electron Configurations 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d To write the electron configuration of an element, you must fill the sublevels in order of increasing energy. If you follow the arrows in either of the two types of mnemonics shown below, you will get correct configurations for most elements. 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 7p 3d 4d 5d 6d 4f 5f Math Tutor Answers 1. a. 1s 2 2s 2 2p 6 3s 2 3p 1 b. 1s 2 2s 2 2p 6 c. 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 5s 2 5p 2 d. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 2. a. [Ne3s 2 3p 2 b. [Kr5s 1 c. [Kr4d 10 5s 2 5p 3 d. [Ar3d 10 4s 2 4p 3 You also need to know how many orbitals are in each sublevel and that each orbital can contain two electrons of opposite spins. As shown in the following table, the sublevels s, p, d, and f have 1, 3, 5, and 7 available orbitals, respectively. SUBLEVEL s p d f No. of orbitals 1 3 5 7 No. of electrons 2 6 10 14 Sample Problem Write the full electron configuration for phosphorus. The atomic number of phosphorus is 15, so a phosphorus atom has 15 protons and electrons. Assign each of the 15 electrons to the appropriate sublevels. The final sublevel can be unfilled and will contain the number of valence electrons. 1s 2 2s 2 2p 6 3s 2 3p 3 2e - + 2e - + 6e - + 2e - + 3e - = 15e - So, the full electron configuration of phosphorus is 1s 2 2s 2 2p 6 3s 2 3p 3. Answers in Appendix E 1. Write full electron configurations for the following elements. a. aluminum c. tin b. neon d. potassium 2. Use noble gas symbols to write shorthand electron configurations for the following elements. a. silicon c. antimony b. rubidium d. arsenic Math Tutor 157 led-46 157 5/9/2011 10:55:27 AM The Periodic Law 157

Chapter summary CHAPTER 5 Summary PREMIUM CONTENT Interactive Review HMDScience.com Review Games Concept Maps SECTION 1 History of the Periodic Table The periodic law states that the physical and chemical properties of the elements are periodic functions of their atomic numbers. The periodic table is an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column. The columns in the periodic table are referred to as groups. KEY TERMS periodic law periodic table lanthanide actinide SECTION 2 Electron Configuration and the Periodic Table The rows in the periodic table are called periods. Many chemical properties of the elements can be explained by the configurations of the elements outermost electrons. The noble gases exhibit unique chemical stability because their highest occupied levels have an octet of electrons, ns 2 np 6 (with the exception of helium, whose stability arises from its highest occupied level being completely filled with two electrons, 1s 2 ). Based on the electron configurations of the elements, the periodic table can be divided into four blocks: the s-block, the p-block, the d-block, and the f- block. KEY TERMS alkali metals alkaline-earth metals transition elements main-group elements halogens SECTION 3 Electron Configuration and Periodic Properties The groups and periods of the periodic table display general trends in the following properties of the elements: electron affinity, electronegativity, ionization energy, atomic radius, and ionic radius. The electrons in an atom that are available to be lost, gained, or shared in the formation of chemical compounds are referred to as valence electrons. In determining the electron configuration of an ion, the order in which electrons are removed from the atom is the reverse of the order given by the atom s electron - configuration notation. KEY TERMS atomic radius ion ionization ionization energy electron affinity cation anion valence electron electronegativity 158 Chapter 5 Review Answers 1. a. Cannizzaro developed a standard method for measuring atomic masses, which allowed chemists to search for periodic trends among elements. b. Mendeleev organized elements according to increasing atomic mass and noticed that similar properties appeared periodically. c. Moseley discovered that nuclear charge (i.e., atomic number), not atomic mass, should be the basis for organizing the periodic table. 2. The physical and chemical properties of the elements are periodic functions of their atomic numbers. 3. Groups of elements exhibit similar chemical and physical properties and behavior. 4. a. Generally, the configurations of the outermost electron shells of elements within the same group are the same. There are a number of exceptions to this rule, however, among the transition elements. b. Their outer shells are completely filled. Untitled-46 158 5/9/2011 10:55:28 158 Chapter 5

titled-46 159 CHAPTER 5 Review SECTION 1 History of the Periodic Table 1. Describe the contributions made by the following scientists to the development of the periodic table: a. Stanislao Cannizzaro b. Dmitri Mendeleev c. Henry Moseley 2. State the periodic law. 3. How is the periodic law demonstrated within the groups of the periodic table? SECTION 2 Electron Configuration and the Periodic Table 4. a. How do the electron configurations within the same group of elements compare? b. Why are the noble gases relatively unreactive? 5. What determines the length of each period in the periodic table? 6. What is the relationship between the electron configuration of an element and the period in which that element appears in the periodic table? 7. a. What information is provided by the specific block location of an element? b. Identify, by number, the groups located within each of the four block areas. 8. a. Which elements are designated as the alkali metals? b. List four of their characteristic properties. 9. a. Which elements are designated as the alkalineearth metals? b. How do their characteristic properties compare with those of the alkali metals? 10. a. Write the group configuration notation for each d-block group. b. How do the group numbers of those groups relate to the number of outer s and d electrons? 5. The length of a period is determined by the total number of electrons that can fill the outer sublevels of the elements of that period. 6. An element s period corresponds to its highest occupied main energy level. 7. a. the type of sublevel being filled in successive elements of that block b. s-block: Groups 1 and 2; p-block: Groups 13 18 (except for He); d-block: Groups 3 12 (except for the f block elements); f-block: those elements of the sixth and seventh periods between Groups 3 and 4 11. What name is sometimes used to refer to the entire set of d-block elements? 12. a. What types of elements make up the p-block? b. How do the properties of the p-block metals compare with those of the metals in the s- and d-blocks? 13. a. Which elements are designated as the halogens? b. List three of their characteristic properties. 14. a. Which elements are metalloids? b. Describe their characteristic properties. 15. Which elements make up the f-block in the periodic table? 16. a. What are the main-group elements? b. What trends can be observed across the various periods within the main-group elements? PRACTICE PROBLEMS 17. Write the noble-gas notation for the electron configuration of each of the following elements, and indicate the period in which each belongs. a. Li c. Cu e. Sn b. O d. Br 18. Without looking at the periodic table, identify the period, block, and group in which the elements with the following electron configurations are located. (Hint: See Sample Problem A.) a. [Ne3s 2 3p 4 b. [Kr4d 10 5s 2 5p 2 c. [Xe4f 14 5d 10 6s 2 6p 5 19. Based on the information given below, give the group, period, block, and identity of each element described. (Hint: See Sample Problem B.) a. [He2s 2 b. [Ne3s 1 c. [Kr5s 2 d. [Ar4s 2 e. [Ar3d 5 4s 1 20. Without looking at the periodic table, write the expected outer electron configuration for each of the following elements. (Hint: See Sample Problem C.) a. Group 7, fourth period b. Group 3, fifth period c. Group 12, sixth period Chapter Review 159 8. a. the Group 1 elements b. The Group 1 elements are extremely 5/9/2011 10:55:29 AM reactive. They react vigorously with water, they are silvery in appearance, and each is soft enough to be cut with a knife. 9. a. the Group 2 elements b. The Group 2 elements are harder, denser, and stronger than the Group 1 elements. They also have higher melting points. Group 2 elements are less reactive than Group 1 elements. 10. a. Group 3: (n 1)d 1 ns 2 Group 4: (n 1)d 2 ns 2 Group 5: (n 1)d 3 ns 2 Group 6: (n 1)d 5 ns 1 for Cr and Mo (n 1)d 4 ns 2 for W and Sg Group 7: (n 1)d 5 ns 2 Group 8: (n 1)d 6 ns 2 Group 9: (n 1)d 7 ns 2 Group 10: (n 1)d 8 ns 2 Group 11: (n 1)d 10 ns 1 Group 12: (n 1)d 10 ns 2 b. The group number equals the sum of the outer s and d electrons. 11. the transition elements 12. a. The p-block consists of nonmetals at the right, metalloids in the middle, and metals at the left. b. The p block metals are generally harder and more dense than the s block metals but softer and less dense than the d block metals. 13. a. the Group 17 elements b. The halogens are the most-reactive nonmetals; they react vigorously with most metals to form salts. They are also among the most electronegative elements. 14. a. B, Si, Ge, As, Sb, Te b. The metalloids are mostly brittle solids with electrical conductivities intermediate between those of metals (good conductors) and nonmetals (nonconductors). 15. the lanthanides and the actinides 16. a. the elements of the s and p-blocks (plus hydrogen and helium) b. decrease in atomic size, increase in ionization energy, increase in electron affinity, decrease in cationic size, decrease in anionic size, increase in electronegativity 17. a. [He2s 1, second period b. [He2s 2 2p 4, second period c. [Ar3d 10 4s 1, fourth period d. [Ar3d 10 4s 2 4p 5, fourth period e. [Kr4d 10 5s 2 5p 2, fifth period 18. a. third period, p-block, Group 16 b. fifth period, p-block, Group 14 c. sixth period, p-block, Group 17 19. a. Group 2, second period, s-block, Be b. Group 1, third period, s-block, Na c. Group 2, fifth period, s-block, Sr d. Group 2, fourth period, s-block, Ca e. Group 6, fourth period, d-block, Cr 20. a. 3d 5 4s 2, b. 4d 1 5s 2, c. 4f 14 5d 10 6s 2 The Periodic Law 159

chapter review 21. a. p-block, third period, Group 13, Al, metal, high reactivity b. p-block, fourth period, Group 18, noble gases, Kr, nonmetal, low reactivity c. d-block, fifth period, Group 11, Ag, metal, low reactivity d. f-block, sixth period, between Groups 3 and 4, Ce, metal, high reactivity 22. a. one half the distance between the nuclei of two bonded identical atoms b. They decrease. c. As electrons are added to s and p sublevels in the same main energy level, the increasing positive charge of the nucleus pulls electrons closer to the nucleus, resulting in decreasing atomic radii. 23. a. They generally increase. b. Down a group, the outer electrons of each element occupy comparable sublevels in successively higher main energy levels farther from the nucleus. 24. a. a charged atom or a charged group of bonded atoms b. any process that results in the formation of an ion c. the energy required to remove one electron from a neutral atom of an element d. the energy required to remove an electron from a 1+ ion 25. a. They increase across a period and decrease down a group. b. Across a period, the increasing nuclear charge attracts electrons in the same energy level more strongly and makes them more difficult to remove. Down a group, the electrons to be removed from each successive element are farther from the nucleus in increasingly higher energy levels and are thus more easily removed. 26. a. the energy taken in or given off when an electron is added to an atom b. Electron affinity values are either negative or positive. A negative sign indicates that energy is given off; a positive sign indicates that energy is taken in. 27. a. A cation is a positive ion, and an anion is a negative ion. 21. Identify the block, period, group, group name (where appropriate), element name, element type, and relative reactivity for the elements with the following electron configurations. (Hint: See Sample Problem D.) a. [Ne3s 2 3p 1 b. [Ar3d 10 4s 2 4p 6 c. [Kr4d 10 5s 1 d. [Xe4f 1 5d 1 6s 2 SECTION 3 Electron Configuration and Periodic Properties 22. a. What is meant by atomic radius? b. What trend is observed among the atomic radii of main-group elements across a period? c. Explain this trend. 23. a. What trend is observed among the atomic radii of main-group elements down a group? b. Explain this trend. 24. Define each of the following terms: a. ion b. ionization c. first ionization energy d. second ionization energy 25. a. How do the first ionization energies of main-group elements vary across a period and down a group? b. Explain the basis for each trend. 26. a. What is electron affinity? b. What signs are associated with electron affinity values, and what is the significance of each sign? 27. a. Distinguish between a cation and an anion. b. How does the size of each compare with the size of the neutral atom from which it is formed? 28. a. What are valence electrons? b. Where are such electrons located? 160 Chapter 5 b. Cations are always smaller than the atoms from which they are formed; anions are always larger. 28. a. Valence electrons are atomic electrons available to be lost, gained, or shared in the formation of chemical compounds. b. Valence electrons are located in an atom s outermost energy level. 29. a. lost, 1 d. gained, 2 b. lost, 2 e. gained, 1 c. lost, 3 f. neither lost nor gained, 0 30. a. Electronegativity is the ability of an atom in a chemical compound to attract 29. For each of the following groups, indicate whether electrons are more likely to be lost or gained in compound formation, and give the number of such electrons typically involved. a. Group 1 d. Group 16 b. Group 2 e. Group 17 c. Group 13 f. Group 18 30. a. What is electronegativity? b. Why is fluorine special in terms of electronegativity? 31. Identify the most- and least-electronegative groups of elements in the periodic table. PRACTICE PROBLEMS 32. Of cesium, Cs, hafnium, Hf, and gold, Au, which element has the smallest atomic radius? Explain your answer in terms of trends in the periodic table. (Hint: See Sample Problem E.) 33. a. Distinguish between the first, second, and third ionization energies of an atom. b. How do the values of successive ionization energies compare? c. Why does this occur? 34. Without looking at the electron affinity table, arrange the following elements in order of decreasing electron affinities: C, O, Li, Na, Rb, and F. 35. a. Without looking at the ionization energy table, arrange the following elements in order of decreasing first ionization energies: Li, O, C, K, Ne, and F. b. Which of the elements listed in (a) would you expect to have the highest second ionization energy? Why? 36. a. Which of the following cations is least likely to form: Sr 2+, Al 3+, K 2+? b. Which of the following anions is least likely to form: I -, Cl -, O 2-? 37. Which element is the most electronegative among C, N, O, Br, and S? Which group does it belong to? (Hint: See Sample Problem G.) 38. The two ions K + and Ca 2+ each have 18 electrons surrounding the nucleus. Which would you expect to have the smaller radius? Why? electrons from other atoms. b. Fluorine is the most electronegative element and is arbitrarily assigned an electronegativity of 4.0. The values for all other elements are assigned in relation to this value. 31. Most-electronegative: Group 17, least-electronegative: Group 1 32. Gold; atomic radii decrease across a period, and gold is farthest to the right in the sixth period, in which all 3 elements are found. 33. a. First ionization energy is the energy required to remove one electron from an Untitled-46 160 5/9/2011 10:55:31 160 Chapter 5

titled-46 161 Mixed Review 39. Without looking at the periodic table, identify the period, block, and group in which each of the following elements is located. a. [Rn7s 1 b. [Ar3d 2 4s 2 c. [Kr4d 10 5s 1 d. [Xe4f 14 5d 9 6s 1 40. a. Which elements are designated as the noble gases? b. What is the most significant property of these elements? 41. Which of the following does not have a noble-gas configuration: Na +, Rb +, O 2-, Br - Ca +, Al 3+, S 2-? 42. a. How many groups are in the periodic table? b. How many periods are in the periodic table? c. Which two blocks of the periodic table make up the main-group elements? 43. Write the noble-gas notation for the electron configuration of each of the following elements, and indicate the period and group in which each belongs. a. Mg b. P c. Sc d. Y 44. Use the periodic table to describe the chemical properties of the following elements: a. fluorine, F b. xenon, Xe c. sodium, Na d. gold, Au 45. For each element listed below, determine the charge of the ion that is most likely to be formed and the identity of the noble gas whose electron configuration is thus achieved. a. Li e. Mg i. Br b. Rb f. Al j. Ba c. O g. P d. F h. S 46. Describe some differences between the s-block metals and the d-block metals. atom, which is neutral. Second ionization energy is the energy required to remove an electron from a 1+ ion. Third ionization energy is the energy required to remove an electron from a 2+ ion. b. Each successive ionization energy is larger than the preceding one. c. Each successive electron must be removed from a particle with a greater positive charge. 34. In order of decreasing electron affinities, the elements are F, O, C, Li, Na, and Rb. 35. a. Ne, F, O, C, Li, K 47. Why do the halogens readily form 1 ions? 48. Identify which trends in the diagrams below describe atomic radius, ionization energy, electron affinity, and electronegativity. a. b. c. 49. The electron configuration of argon differs from those of chlorine and potassium by one electron each. Compare the reactivity of these three elements. Chapter Review 161 b. Li and K would have the highest second ionization energies, because in both cases 5/9/2011 10:55:32 AM the second electron must come from a completely filled noble-gas electron configuration. Of the two, Li would have the higher second ionization energy, because the Li + ion is smaller than the K + ion. 36. a. K 2+, b. O 2 37. O; Group 16 38. Ca 2+ ; because of the greater attraction that its 20 protons exert on the 18 electrons 39. a. seventh period, s-block, Group 1 b. fourth period, d-block, Group 4 c. fifth period, d-block, Group 11 d. sixth period, d-block, Group 10 40. a. the Group 18 elements b. They are very unreactive. 41. Ca + 42. a. 18; b. 7; c. the s and p-blocks 43. a. [Ne3s 2, third period, Group 2 b. [Ne3s 2 3p 3, third period, Group 15 c. [Ar3d 1 4s 2, fourth period, Group 3 d. [Kr4d 1 5s 2, fifth period, Group 3 44. a. It is a highly reactive nonmetal in Group 17, Period 2. It needs one electron to achieve a noble gas configuration, which explains its high electron affinity. It has a high ionization energy, and thus the formation of positive ions is not likely. Fluorine is the most electronegative element. b. It is a nonmetal of low reactivity in Group 18, Period 5. It has a filled outer energy level, so there is little tendency to lose, gain, or share electrons. It has the highest ionization energy of the elements in the fifth period. c. It is a highly reactive metal in Group 1, Period 3. It has a low first ionization energy, because losing an electron to form Na + gives it a noble gas configuration. It has a low electron affinity and a low electronegativity. d. It is a metal of fairly low reactivity in Group 11, Period 6. Its ionization energy is relatively low, so it forms positive ions. 45. a. 1 +, He f. 3+, Ne b. 1+, Kr g. 3, Ar c. 2, Ne h. 2, Ar d. 1, Ne i. 1, Kr e. 2+, Ne j. 2+, Xe 46. The d block metals are harder, denser, and less reactive, and except for mercury, they have higher melting points than s block metals do. 47. In doing so, they achieve a noble-gas configuration. 48. ionization energy, electron affinity, electronegativity: diagram a; atomic radius: diagram b 49. Chlorine and potassium are reactive elements, whereas argon is unreactive. Chlorine readily gains one electron to achieve the configuration of argon, and potassium readily loses one electron to achieve the argon configuration. The Periodic Law 161

chapter review 50. 4 Λ 9 12 γ 13 β 23 ξ 51. 5 52. 8 53. From left to right, the elements are organized according to their atomic numbers. Elements are organized into families or columns according to their properties. The elements exhibit increasingly nonmetallic behavior from left to right across the table. Their physical states suggest that melting points and boiling points also increase from left to right. 54. Na is a liquid at 97.8 C. K is a liquid at 63.25 C. Rb is a liquid at 38.89 C. Cs is a liquid at 28.5 C. Fr is a liquid at 27 C. Hg is a liquid at 38.8 C. Ga is a liquid at 29.77 C. N as N 2 is a gas at 195.8 C. P is a liquid at 44.1 C. O as O 2 is a gas at 182.962 C. F as F 2 is a gas at 188.14 C. Cl as Cl 2 is a gas at 34.6 C. Br as Br 2 is a gas at 58.78 C. 55. a. Cr 3+, violet or green; Cr 6+, yellow or orange b. rubies c. copper, blue; cadmium, yellow; cobalt, pink; zinc, white; nickel, green d. Gold, silver, platinum, palladium, iridium, rhodium, ruthenium, osmium; they have low reactivity. 56. Possible responses include Friedrich Dorn s discovery of radon in 1900, Moseley s determination of the nuclear charge (i.e., atomic number) of each element by 1913, the discovery of the lanthanides by 1913, the synthesis in 1940 of the first element beyond uranium (neptunium) in the periodic table by Philip Abelson and Edwin McMillan, the creation of plutonium in 1941 by Seaborg, and the synthesis of element 109 in Germany in 1982. CRITICAL THINKING As a member on the newly-inhabited space station Alpha, you are given the task of organizing information on newly discovered elements as it comes in from the laboratory. To date, five elements have been discovered and have been assigned names and symbols from the Greek alphabet. An analysis of the new elements has yielded the following data: Element name Epsilon ε Beta β Gamma γ Delta Lambda Λ Atomic no. Atomic mass Properties 23 47.33 nonmetal, very reactive, produces a salt when combined with a metal, gaseous state 13 27.01 metal, very reactive, soft solid, low melting point 12 25.35 nonmetal, gaseous element, extremely unreactive 4 7.98 nonmetal, very abundant, forms compounds with most other elements 9 16.17 metal, solid state, good conductor, high luster, hard and dense 50. Applying Models Create a periodic table based on the properties of the five new elements. 51. Predicting Outcomes Using your newly created periodic table, predict the atomic number of an element with an atomic mass of 11.29 that has nonmetallic properties and is very reactive. 52. Predicting Outcomes Predict the atomic number of an element having an atomic mass of 15.02 that exhibits metallic properties but is softer than lambda and harder than beta. 53. Analyzing Information Analyze your periodic table for trends, and describe those trends. 162 Chapter 5 57. Reports should mention Seaborg s various roles in the syntheses of curium (element 96) in 1944, californium (element 98) in 1950, mendelevium (el ement 101) in 1955, and nobelium (element 102) in 1958. 58. Student examples may vary, but should include isolated trends arranged in a periodic manner by columns and rows, as in the periodic table of the elements. USING THE HANDBOOK 54. Review the boiling point and melting point data in the tables of the Elements Handbook (Appendix A). Make a list of the elements that exist as liquids or gases at the boiling point of water, 100 C. 55. Because transition metals have vacant d orbitals, they form a greater variety of colored compounds than do the metals of Groups 1 and 2. Review the section of the Elements Handbook (Appendix A) on transition metals, and answer the following: a. What colors are exhibited by chromium in its common oxidation states? b. What gems contain chromium impurities? c. What colors are often associated with the following metal ions: copper, cadmium, cobalt, zinc, and nickel? d. What transition elements are considered noble metals? What are the characteristics of a noble metal? RESEARCH AND WRITING 56. Prepare a report tracing the evolution of the current periodic table since 1900. Cite the chemists involved and their major contributions. 57. Write a report describing the contributions of Glenn Seaborg toward the discovery of many of the actinide elements. ALTERNATIVE ASSESSMENT 58. Construct your own periodic table or obtain a poster that shows related objects, such as fruits or vegetables, in periodic arrangement. Describe the organization of the table and the trends it illustrates. Use this table to make predictions about your subject matter. Untitled-46 162 5/9/2011 10:55:33 162 Chapter 5