1. Periodicity

2. Classification of the Elements u OBJECTIVES: Explain why you can infer the properties of an element based on those of other elements in the periodic table.

3. Classification of the Elements u OBJECTIVES: Use electron configurations to classify elements as noble gases, main group elements, transition metals, or inner transition metals.

4. Periodic Table Revisited u Russian scientist Dmitri Mendeleev taught chemistry in terms of properties. u Mid 1800’s - molar masses of elements were known. u Wrote down the elements in order of increasing mass. u Found a pattern of repeating properties.

5. Mendeleev’s Table u Grouped elements in columns by similar properties in order of increasing atomic mass. u Found some inconsistencies - felt that the properties were more important than the mass, so switched order. u Also found some gaps. u Must be undiscovered elements. u Predicted their properties before they were found.

6. The modern table u Elements are still grouped by properties. u Similar properties are in the same column. u Order is by increasing atomic number. u Added a column of elements Mendeleev didn’t know about. u The noble gases weren’t found because they didn’t react with anything.

7. u Horizontal rows are called periods u There are 7 periods

8. Vertical columns called groups Elements are placed in columns by similar properties Also called families

9. 1A 2A3A4A5A6A 7A 8A 0 u The elements in the A groups are called the representative elements outer s or p filling

10. The group B are called the transition elements u These are called the inner transition elements, and they belong here

11. u Group 1A are the alkali metals u Group 2A are the alkaline earth metals

12. u Group 7A is called the Halogens u Group 8A are the noble gases

13. Why? u The part of the atom another atom sees is the electron cloud. u More importantly the outside orbitals. u The orbitals fill up in a regular pattern. u The outside orbital electron configuration repeats. u The properties of atoms repeat.

14. 1s11s1 1s 2 2s 1 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6 s 2 4f 14 5d 10 6p 6 7s 1 H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87

15. He 2 Ne 10 Ar 18 Kr 36 Xe 54 Rn 86 1s21s2 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 2 3p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6

16. u Alkali metals all end in s 1 u Alkaline earth metals all end in s 2 u really should include He, but it fits better later. u He has the properties of the noble gases. s2s2 s1s1 S- block

17. Transition Metals -d block d1d1 d2d2 d3d3 s1d5s1d5 d5d5 d6d6 d7d7 d8d8 s 1 d 10 d 10

18. The P-block p1p1 p2p2 p3p3 p4p4 p5p5 p6p6

19. F - block u inner transition elements

20. u Each row (or period) is the energy level for s and p orbitals. 12345671234567

21. u d orbitals fill up after previous energy level, so first d is 3d even though it’s in row 4. 12345671234567 3d

22. u f orbitals start filling at 4f 12345671234567 4f 5f

23. Writing electron configurations the easy way

24. Electron Configurations repeat u The shape of the periodic table is a representation of this repetition. u When we get to the end of the column the outermost energy level is full. u This is the basis for our shorthand.

25. The Shorthand u Write symbol of the noble gas before the element, in [ ]. u Then, the rest of the electrons. u Aluminum’s full configuration: 1s 2 2s 2 2p 6 3s 2 3p 1 u previous noble gas Ne is: 1s 2 2s 2 2p 6 u so, Al is: [Ne] 3s 2 3p 1

26. More examples u Ge = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 2 Thus, Ge = [Ar] 4s 2 3d 10 4p 2 u Hf = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 2 Thus, Hf = [Xe]6s 2 4f 14 5d 2

27. The Shorthand Again Sn- 50 electrons The noble gas before it is Kr [ Kr ] Takes care of 36 Next 5s 2 5s 2 Then 4d 10 4d 10 Finally 5p 2 5p 2

28. Periodic Trends u OBJECTIVES: Interpret group trends in atomic radii, ionic radii, ionization energies, m.p., b.p., electronegativity and chemical properties

29. Trends in Atomic Size u First problem: Where do you start measuring from? u The electron cloud doesn’t have a definite edge. u They get around this by measuring more than 1 atom at a time.

30. Atomic Size u Atomic Radius = half the distance between two nuclei of a diatomic molecule. } Radius

31. Trends in Atomic Size u Influenced by three factors: 1. Energy Level Higher energy level is further away. 2. Charge on nucleus More charge pulls electrons in closer. u 3. Shielding effect e e repulsion

32. Group trends u As we go down a group... u each atom has another energy level, u so the atoms get bigger. H Li Na K Rb

33. Periodic Trends u As you go across a period, the radius gets smaller. u Electrons are in same energy level. u More nuclear charge. u Outermost electrons are closer. NaMgAlSiPSClAr

34. Overall Atomic Number Atomic Radius (nm) H Li Ne Ar 10 Na K Kr Rb

35. Trends in Ionization Energy u The amount of energy required to completely remove a mole of electrons from a mole of gaseous atoms. u Removing an electron makes a +1 ion. u The energy required to remove (1 mole of) the first electron is called the first ionization energy.

36. Ionization Energy u The second ionization energy is the energy required to remove (1 mole of) the second electron(s). u Always greater than first IE. u The third IE is the energy required to remove a third electron. u Greater than 1st or 2nd IE.

37. SymbolFirstSecond Third H He Li Be B C N O F Ne 1312 2731 520 900 800 1086 1402 1314 1681 2080 5247 7297 1757 2430 2352 2857 3391 3375 3963 11810 14840 3569 4619 4577 5301 6045 6276

38. SymbolFirstSecond Third H He Li Be B C N O F Ne 1312 2731 520 900 800 1086 1402 1314 1681 2080 5247 7297 1757 2430 2352 2857 3391 3375 3963 11810 14840 3569 4619 4577 5301 6045 6276

39. What determines IE u The greater the nuclear charge, the greater IE. u Greater distance from nucleus decreases IE u Filled and half-filled orbitals have lower energy, so achieving them is easier, lower IE. u Shielding effect

40. Shielding u The electron on the outermost energy level has to look through all the other energy levels to see the nucleus. u Second electron has same shielding, if it is in the same period

41. Group trends u As you go down a group, first IE decreases because... u The electron is further away. u More shielding.

42. Periodic trends u All the atoms in the same period have the same energy level. u Same shielding. u But, increasing nuclear charge u So IE generally increases from left to right. u Exceptions at full and 1/2 full orbitals.

43. First Ionization energy Atomic number He u He has a greater IE than H. u same shielding u greater nuclear charge H

44. First Ionization energy Atomic number H He Li has lower IE than H more shielding further away l outweighs greater nuclear charge Li

45. First Ionization energy Atomic number H He Be has higher IE than Li same shielding l greater nuclear charge Li Be

46. First Ionization energy Atomic number H He B has lower IE than Be same shielding greater nuclear charge l p orbital is slightly more diffuse and its electron easier to remove Li Be B

47. First Ionization energy Atomic number H He Li Be B C

48. First Ionization energy Atomic number H He Li Be B C N

49. First Ionization energy Atomic number H He Li Be B C N O u Breaks the pattern, because removing an electron leaves 1/2 filled p orbital

50. First Ionization energy Atomic number H He Li Be B C N O F

51. First Ionization energy Atomic number H He Li Be B C N O F Ne u Ne has a lower IE than He u Both are full, u Ne has more shielding u Greater distance

52. First Ionization energy Atomic number H He Li Be B C N O F Ne Na has a lower IE than Li Both are s 1 Na has more shielding l Greater distance Na

53. First Ionization energy Atomic number

54. Driving Force u Full Energy Levels require lots of energy to remove their electrons. u Noble Gases have full orbitals. u Atoms behave in ways to achieve noble gas configuration.

55. 2nd Ionization Energy u For elements that reach a filled or half-filled orbital by removing 2 electrons, 2nd IE is lower than expected. u True for s 2 u Alkaline earth metals form 2+ ions.

56. 3rd IE u Using the same logic s 2 p 1 atoms have an low 3rd IE. u Atoms in the aluminum family form 3+ ions. u 2nd IE and 3rd IE are always higher than 1st IE!!!

57. Trends in Electron Affinity u The energy change associated with adding an electron to a gaseous atom. u Easiest to add to group 7A. u Gets them to full energy level. u Increase from left to right: atoms become smaller, with greater nuclear charge. u Decrease as we go down a group.

58. Trends in Ionic Size u Cations form by losing electrons. u Cations are smaller that the atom they come from. u Metals form cations. u Cations of representative elements have noble gas configuration.

59. Ionic size u Anions form by gaining electrons. u Anions are bigger that the atom they come from. u Nonmetals form anions. u Anions of representative elements have noble gas configuration.

60. Configuration of Ions u Ions always have noble gas configuration. u Na is: 1s 2 2s 2 2p 6 3s 1 u Forms a 1+ ion: 1s 2 2s 2 2p 6 u Same configuration as neon. u Metals form ions with the configuration of the noble gas before them - they lose electrons.

61. Configuration of Ions u Non-metals form ions by gaining electrons to achieve noble gas configuration. u They end up with the configuration of the noble gas after them.

62. Group trends u Adding energy level u Ions get bigger as you go down. Li 1+ Na 1+ K 1+ Rb 1+ Cs 1+

63. Periodic Trends u Across the period, nuclear charge increases so they get smaller. u Energy level changes between anions and cations. Li 1+ Be 2+ B 3+ C 4+ N 3- O 2- F 1-

64. Size of Isoelectronic ions u Iso- means the same u Iso electronic ions have the same # of electrons u Al 3+ Mg 2+ Na 1+ Ne F 1- O 2- and N 3- u all have 10 electrons u all have the configuration: 1s 2 2s 2 2p 6

65. Size of Isoelectronic ions u Positive ions that have more protons would be smaller. Al 3+ Mg 2+ Na 1+ Ne F 1- O 2- N 3-

66. Electronegativity u The tendency for an atom to attract electrons to itself when it is chemically combined with another element. u High electronegativity means it pulls the electron toward it. u Atoms with large negative electron affinity have larger electronegativity.

67. Group Trend u The further down a group, the farther the electron is away, and the more electrons an atom has. u More willing to share. u Low electronegativity.

68. Periodic Trend u Metals are at the left of the table. u They let their electrons go easily u Low electronegativity u At the right end are the nonmetals. u They want more electrons. u Try to take them away from others u High electronegativity.

69. Ionization energy, Electronegativity, and Electron Affinity INCREASE

70. Atomic size increases, shielding constant Ionic size increases