1. Periodic Properties of the Elements Chapter 7 Periodic Properties of the Elements Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten John D. Bookstaver St. Charles Community College St. Peters, MO  2006, Prentice Hall, Inc.

2. Periodic Properties of the Elements SIZE OF ATOMS AND IONS IONIZATION ENERGY ELECTRON AFFINITY GROUP TRENDS HW - ONLINE QUESTIONS PRACTICE 1 2, 9, 13, 14, 21, 23, 27, 28, 29, 31, 33, 37, 41, 43, 45, 57, 60, 63, 67, 73, 77 MONDAY BRING EQUATIONS BOOK

3. Periodic Properties of the Elements Homework New Atomic Theory Exams Questions is a file in the John Bowne site. Has similar questions to AP test. Practice them. The more you do them the better you’ll do in the test. Test Unit 6 and 7 Monday 26

4. Periodic Properties of the Elements The periodic table can be used as a guide for electron configurations. The period number is the value of n. Groups 1 and 2 have the s-orbital filled. Groups 13 - 18 have the p-orbitals filled. Groups 3 - 12 have the d-orbitals filled. The lanthanides and actinides have the f-orbital filled. Electron Configurations and the Periodic Table

5. Periodic Properties of the Elements Periodic Table and electron configuration We fill orbitals in increasing order of energy. Different blocks on the periodic table, then correspond to different types of orbitals.

6. Periodic Properties of the Elements Development of Periodic Table Elements in the same group generally have similar chemical properties. Properties are not identical, however.

7. Periodic Properties of the Elements Development of Periodic Table Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped in 1869.

8. Periodic Properties of the Elements PERIODIC ARRANGEMENT Mendeleev arranged the elements according to their atomic masses. In 1913 Henry Moseley developed the concept of ATOMIC NUMBERS which lead to the current arrangement according to their ATOMIC NUMBER.

9. Periodic Properties of the Elements EffectiveEffective Nuclear Charge In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons. The nuclear charge that an electron experiences depends on both factors.

10. Periodic Properties of the Elements Effective Nuclear Charge The effective nuclear charge, Z eff, is found this way: Z eff = Z − S where Z is the atomic number and S is a screening constant, usually close to the number of inner electrons.

11. Periodic Properties of the Elements Effective Nuclear Charge The effective nuclear charge is smaller than the actual nuclear charge because the effective nuclear charge Z eff accounts for the repulsion of the electron by the other electrons in the atom. Z eff < Z

12. Periodic Properties of the Elements Core electrons shield or screen the outer electrons from the attraction of the nucleus.

13. Periodic Properties of the Elements Sizes of Atoms The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.

14. Periodic Properties of the Elements Sizes of Atoms Bonding atomic radius tends to… …decrease from left to right across a row due to increasing Z eff. …increase from top to bottom of a column due to increasing value of n

15. Periodic Properties of the Elements Examples – Place each group of elements in order of increasing atomic radius: 1.S, Al, Cl, Mg, Ar, Na 2.K, Li, Cs, Na, H 3.Ca, As, F, Rb, O, K, S, Ga

16. Periodic Properties of the Elements Examples – Place each group of elements in order of increasing atomic radius: 1.S, Al, Cl, Mg, Ar, Na Ar < Cl < S < Al < Mg < Na 2.K, Li, Cs, Na, H H < Li < Na < K < Cs 3.Ca, F, As, Rb, O, K, S, Ga F < O < S < As < Ga < Ca < K < Rb

17. Periodic Properties of the Elements Electron Configurations of Ions Cations: electrons removed from orbital with highest principle quantum number, n, first: Li (1s 2 2s 1 )  Li + (1s 2 ) Fe ([Ar]3d 6 4s 2 )  Fe 3+ ([Ar]3d 5 ) Anions: electrons added to the orbital with highest n: F (1s 2 2s 2 2p 5 )  F  (1s 2 2s 2 2p 6 )

18. Periodic Properties of the Elements Write electron configurations for the following ions: 1.Al 3+ 2.S 2- 3.Li + 4.Br - 5.Fe 2+ 6.Fe 3+

19. Periodic Properties of the Elements Write electron configurations for the following ions: 1.Al 3+ 1s 2 2s 2 2p 6 2.S 2- [Ne]3s 2 3p 6 3.Li + 1s 2 4.Br - [Ar]4s 2 3d 10 4p 6 5.Fe 2+ [Ar]3d 6 6.Fe 3+ [Ar]3d 5

20. Periodic Properties of the Elements Sizes of Ions Ionic size depends upon:  Nuclear charge.  Number of electrons.  Orbitals in which electrons reside.

21. Periodic Properties of the Elements Sizes of Ions Cations are smaller than their parent atoms.  The outermost electron is removed and repulsions are reduced.

22. Periodic Properties of the Elements Sizes of Ions Anions are larger than their parent atoms.  Electrons are added and repulsions are increased.

23. Periodic Properties of the Elements Sizes of Ions Ions increase in size as you go down a column.  Due to increasing value of n.

24. Periodic Properties of the Elements DO NOW Write the formula for Potassium Permanganate and Sodium hypochlorite Write the electronic configuration of both cations. Indicate to what element are each isoelectronic. Which one has the largest radius?

25. Periodic Properties of the Elements Sizes of Ions In an isoelectronic series, ions have the same number of electrons. Ionic size decreases with an increasing nuclear charge.

26. Periodic Properties of the Elements Examples – Choose the larger species in each case: 1.Na or Na + 2.Br or Br - 3.N or N 3- 4.O - or O 2- 5.Mg 2+ or Sr 2+ 6.Mg 2+ or O 2- 7.Fe 2+ or Fe 3+

27. Periodic Properties of the Elements Examples – Choose the larger species in each case: 1.Na or Na + 2.Br or Br - 3.N or N 3- 4.O - or O 2- 5.Mg 2+ or Sr 2+ 6.Mg 2+ or O 2- 7.Fe 2+ or Fe 3+

28. Periodic Properties of the Elements Ionization Energy Amount of energy required to remove an electron from the ground state of a gaseous atom or ion.  First ionization energy is that energy required to remove first electron.  Second ionization energy is that energy required to remove second electron, etc.

29. Periodic Properties of the Elements IONIZATIONIONIZATION ENERGY THE GREATER THE IONIZATION ENERGY THE MORE DIFFICULT IT IS TO REMOVE AN ELECTRON! It requires more energy to remove each successive electron. When all valence electrons have been removed, the ionization energy takes a quantum leap (removing core electrons require huge amounts of energy).

30. Periodic Properties of the Elements

31. Periodic Properties of the Elements Trends in First Ionization Energies As one goes down a column, less energy is required to remove the first electron.  For atoms in the same group, Z eff is essentially the same, but the valence electrons are farther from the nucleus.

32. Periodic Properties of the Elements Trends in First Ionization Energies Generally, as one goes across a row, it gets harder to remove an electron.  As you go from left to right, Z eff increases.

33. Periodic Properties of the Elements Irregularities N 1402 [He] 2s 2 2p 3 2p 3 this configuration is more stable because half filled orbitals minimize repulsions between electrons O 1314 [He] 2s 2 2p 4 The decrease is due to repulsions of paired electrons in 2p 4

34. Periodic Properties of the Elements Trend across a period IE icreases The energy needed to remove an electron increases because the effective nuclear charge increases and the atomic radius decreases. Irregularities Be 899 B 801 [He] 2s 2 [He] 2s 2 2p 1

35. Periodic Properties of the Elements Trends in First Ionization Energies However, there are two apparent discontinuities in this trend.

36. Periodic Properties of the Elements Trends in First Ionization Energies The first occurs between Groups IIA and IIIA. Electron removed from p-orbital rather than s- orbital  Electron farther from nucleus  Small amount of repulsion by s electrons.

37. Periodic Properties of the Elements Trends in First Ionization Energies The second occurs between Groups VA and VIA.  Electron removed comes from doubly occupied orbital.  Repulsion from other electron in orbital helps in its removal.

38. Periodic Properties of the Elements Examples – Put each set in order of increasing first ionization energy: 1.P, Cl, Al, Na, S, Mg 2.Ca, Be, Ba, Mg, Sr 3.Ca, F, As, Rb, O, K, S, Ga

39. Periodic Properties of the Elements Examples – Put each set in order of increasing first ionization energy: 1.P, Cl, Al, Na, S, Mg 2.Ca, Be, Ba, Mg, Sr 3.Ca, F, As, Rb, O, K, S, Ga 1. Na < Al < Mg < S < P < Cl 2.Ba < Sr < Ca < Mg < Be 3.Rb < K < Ga < Ca < As < S < O < F

40. Periodic Properties of the Elements ElectronElectron Affinity It measures the ease with which and atom gains an electron

41. Periodic Properties of the Elements Electron Affinity Energy change accompanying addition of electron to gaseous atom: Cl + e −  Cl − For most atoms energy is released when an electron is added. The change is exothermic. The more exothermic the most stable the product.

42. Periodic Properties of the Elements Trends in Electron Affinity In general, electron affinity becomes more exothermic as you go from left to right across a row.

43. Periodic Properties of the Elements Trends in Electron Affinity There are again, however, two discontinuities in this trend.

44. Periodic Properties of the Elements Trends in Electron Affinity The first occurs between Groups IA and IIA.  Added electron must go in p-orbital, not s- orbital.  Electron is farther from nucleus and feels repulsion from s-electrons.

45. Periodic Properties of the Elements Trends in Electron Affinity The second occurs between Groups IVA and VA.  Group VA has no empty orbitals.  Extra electron must go into occupied orbital, creating repulsion.

46. Periodic Properties of the Elements Do now Find the Oxidation Number for Oxygen in KO2 K2O2 KO2 Name the compounds

47. Periodic Properties of the Elements Electronegativity: The ability of an atom in a molecule to attract electrons to itself. Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F). Values are calculated from ionization energies and electron affinities. Electronegativity increases across a period and Up a group.

48. Periodic Properties of the Elements Electronegativity

49. Periodic Properties of the Elements Examples – put each set in order by increasing electronegativity: 1.Na, Li, Rb, K, Fr 2.Cl, Ca, F, P, Mg, S, K

50. Periodic Properties of the Elements Examples – put each set in order by increasing electronegativity: 1.Na, Li, Rb, K, Fr 2.Cl, Ca, F, P, Mg, S, K 1.Fr < Rb < K < Na < Li 2.K < Ca < Mg < P < S < Cl < F

51. Periodic Properties of the Elements Properties of Metal, Nonmetals, and Metalloids

52. Periodic Properties of the Elements Metals versus Nonmetals Differences between metals and nonmetals tend to revolve around these properties.

53. Periodic Properties of the Elements Metals versus Nonmetals Metals tend to form cations. Nonmetals tend to form anions.

54. Periodic Properties of the Elements Metals Tend to be lustrous, malleable, ductile, and good conductors of heat and electricity.

55. Periodic Properties of the Elements MetallicMetallic Trend Metals react by losing electrons and forming cations. Metallic character increases with the easiness of losing electrons. The trend is like the ionization energy. Decreases across a period and increases down a group.

56. Periodic Properties of the Elements Metals Compounds formed between metals and nonmetals tend to be ionic. Metal oxides tend to be basic.

57. Periodic Properties of the Elements Nonmetals Dull, brittle substances that are poor conductors of heat and electricity. Tend to gain electrons in reactions with metals to acquire noble gas configuration.

58. Periodic Properties of the Elements Nonmetals Substances containing only nonmetals are molecular compounds. Most nonmetal oxides are acidic.

59. Periodic Properties of the Elements Metalloids Have some characteristics of metals, some of nonmetals. For instance, silicon looks shiny, but is brittle and fairly poor conductor.

60. Periodic Properties of the Elements Group Trends

61. Periodic Properties of the Elements Metals Most metal oxides are basic: (NOT A REDOX RX!!!) Metal oxide + water  metal hydroxide Na 2 O(s) + H 2 O(l)  2NaOH(aq) Nonmetals Nonmetals are more diverse in their behavior than metals. When nonmetals react with metals, nonmetals tend to gain electrons (reduced): metal + nonmetal  salt 2Al(s) + 3Br 2 (l)  2AlBr 3 (s)

62. Periodic Properties of the Elements Nonmetals Most nonmetal oxides are acidic: nonmetal oxide + water  acid (NON A REDOX RX) P 4 O 10 (s) + H 2 O(l)  4H 3 PO 4 (aq) Metalloids Metalloids have properties that are intermediate between metals and nonmetals. Example: Si has a metallic luster but it is brittle. Metalloids have found application in the semiconductor industry.

63. Periodic Properties of the Elements Alkali Metals Soft, metallic solids. Name comes from Arabic word for ashes.

64. Periodic Properties of the Elements moviealkalimetals.mp4 moviealkalimetals.mp4 Alkali Metals Found only as compounds in nature. Very reactive!!! Have low densities and melting points. Also have low ionization energies.

65. Periodic Properties of the Elements Alkali Metals Their reactions with water are famously exothermic. They react vigorously with water and have to be stored in mineral oil or kerosene. 2M (s) + 2 H 2 O (l)  2MOH (aq) + H 2 (g)

66. Periodic Properties of the Elements Alkali Metals Reactions Alkali metals combine directly with most nonmetals Alkali metals (except Li) react with oxygen to form peroxides. (O 2 2- remember ON -1 ) K, Rb, and Cs also form superoxides: K + O 2  KO 2 (ON -1/2!!) Alkali metals produce different oxides when reacting with O 2 : 4Li(s) + O 2 (g)  2Li 2 O(s)(oxide) 2Na(s) + O 2 (g)  Na 2 O 2 (s)(peroxide) K(s) + O 2 (g)  KO 2 (s)(superoxide)

67. Periodic Properties of the Elements With Hydrogen they form HYDRIDES (H - ) LiH, KH With Sulfur they form Sulfides

68. Periodic Properties of the Elements Na line (589 nm): 3p  3s transition Li line: 2p  2s transition K line: 4p  4s transition Group Trends for the Active Metals Group 1: The Alkali Metals Produce bright colors when placed in flame.

69. Periodic Properties of the Elements Alkaline Earth Metals Have higher densities and melting points than alkali metals. Have low ionization energies, but not as low as alkali metals.

70. Periodic Properties of the Elements Alkaline Earth Metals Be does not react with water, Mg reacts only with steam, but others react readily with water. Reactivity tends to increase as go down group. They are less reactive than Alkali metals.

71. Periodic Properties of the Elements Flame Test The flame test is used to visually determine the identity of an unknown metal or metalloid ion based on the characteristic color the salt turns the flame of a bunsen burner. The heat of the flame converts the metal ions into atoms which become excited and emit visible light. The characteristic emission spectra can be used to differentiate between some elements.

72. Periodic Properties of the Elements Flame Test Colors Li + Deep red (crimson) Na + Yellow-orange K+ Violet - lilac Ca 2+ Orange-red Sr 2+ Red Ba 2+ Pale Green Cu 2+ Green

73. Periodic Properties of the Elements Group 6A Oxygen, sulfur, and selenium are nonmetals. Tellurium is a metalloid. The radioactive polonium is a metal.

74. Periodic Properties of the Elements Oxygen Two allotropes: O2O2  O 3, ozone Three anions:  O 2−, oxide  O 2 2−, peroxide  O 2 1−, superoxide Tends to take electrons from other elements (oxidation)

75. Periodic Properties of the Elements Sulfur Weaker oxidizing agent than oxygen. Most stable allotrope is S 8, a ringed molecule.

76. Periodic Properties of the Elements Group VIIA: Halogens Prototypical nonmetals Name comes from the Greek halos and gennao: “salt formers”

77. Periodic Properties of the Elements Group VIIA: HalogensHalogens Large, negative electron affinities  Therefore, tend to oxidize other elements easily React directly with metals to form metal halides Chlorine added to water supplies to serve as disinfectant

78. Periodic Properties of the Elements Group 17: The Halogens The chemistry of the halogens is dominated by gaining an electron to form an anion: X 2 + 2e -  2X -. Fluorine is one of the most reactive substances known: 2F 2 (g) + 2H 2 O(l)  4HF(aq) + O 2 (g)  H = -758.7 kJ. SiO 2 (s) + 2F 2 (g)  SiF 4 (g) + O 2 (g)  H = -704.0 kJ Why is it so reactive? All halogens consist of diatomic molecules, X 2.

79. Periodic Properties of the Elements Chlorine is the most industrially useful halogen. It is produced by the electrolysis of brine (NaCl): 2NaCl(aq) + 2H 2 O(l)  2NaOH(aq) + H 2 (g) + Cl 2 (g). The reaction between chorine and water produces hypochlorous acid (HOCl) which disinfects pool water: Cl 2 (g) + H 2 O(l)  HCl(aq) + HOCl(aq). Halogens react directly with metals to form ionic halides H 2 (g) + X 2  2 HX (g) Hydrogen compounds of the halogens are all strong acids with the exception of HF. LEARN THESE REACTIONS!!!

80. Periodic Properties of the Elements Group VIIIA: Noble Gases Astronomical ionization energies Positive electron affinities  Therefore, relatively unreactive Monatomic gases

81. Periodic Properties of the Elements Group VIIIA: Noble Gases Xe forms three compounds:  XeF 2  XeF 4 (at right)  XeF 6 Kr forms only one stable compound:  KrF 2 The unstable HArF was synthesized in 2000.