1. The Periodic Table

2. History u Russian scientist Dmitri Mendeleev taught chemistry u Mid 1800 - molar masses of elements were known. u Wrote down the elements in order of increasing mass. u Found a pattern of repeating properties.

3. 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 Found some gaps. u Must be undiscovered elements. u Predicted their properties before they were found.

4. The modern table u Elements are still grouped by properties. u Similar properties are in the same column. u Order is in 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.

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

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

7. 1A 2A3A4A5A6A 7A 8A 0 u The elements in the A groups are called the representative elements

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

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

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

11. Why “representative”? u The orbitals fill up in a regular pattern. u The outside orbital electron configuration repeats. u The properties of atoms repeat.

12. 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 6s 2 4f 14 5d 10 6p 6 7s 1 H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87

13. 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

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

15. Transition Metals -d block d1d1 d2d2 d3d3 d4d4 d5d5 d6d6 d7d7 d8d8 d9d9 d 10

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

17. F - block u inner transition elements

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

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

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

21. Writing Electron configurations the easy way Yes there is a shorthand

22. 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 period the outermost energy level is full. u This is the basis for our shorthand.

23. The Shorthand u Write the symbol of the noble gas before the element. u Then the rest of the electrons. u Aluminum - full configuration. u 1s 2 2s 2 2p 6 3s 2 3p 1 u Ne is 1s 2 2s 2 2p 6 u so Al is [Ne] 3s 2 3p 1

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

25. 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

26. Atomic Size u First problem where do you start measuring. u The electron cloud doesn’t have a definite edge. u They get around this by measuring more than 1 atom at a time.

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

28. Trends in Atomic Size Influenced by two factors. 1) Energy Level  Higher energy level is further away. 2) Charge on nucleus (#protons)  More charge pulls electrons in closer.

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

30. Periodic Trends u As you go across a period the radius gets smaller. u Same energy level. u More nuclear charge. u Outermost electrons are closer. NaMgAlSiPSClAr

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

32. Ionization Energy u The amount of energy required to completely remove an electron from a gaseous atom. u Removing one electron makes a +1 ion. u The energy required is called the first ionization energy.

33. u The second ionization energy is the energy required to remove the second electron. u Always greater than first IE. u The third IE is the energy required to remove a third electron. u Greater than 1st & 2nd IE.

34. 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 11 810 14840 3569 4619 4577 5301 6045 6276

35. 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 11 810 14840 3569 4619 4577 5301 6045 6276

36. What determines IE o The greater the nuclear charge the greater IE. o The greater the distance from nucleus, the lower IE o Shielding - lowers the IE

37. Shielding u This outermost electron is shielded from the attractive nuclear forces the inner electrons

38. Group trends u As you go down a group first IE decreases because the electron is further away (more energy levels) and there is more shielding.

39. Periodic trends All the atoms in the same period have the same energy level and same shielding, BUT there is increasing nuclear charge SO IE generally increases from left to right.

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

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

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

43. First Ionization energy Atomic number H He l B has lower IE than Be  same shielding, greater nuclear charge  By removing an electron we make s orbital half filled Li Be B

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

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

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

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

48. 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

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

50. First Ionization energy Atomic number

51. Driving Force u Full Energy Levels are very low energy. u Noble Gases have full orbitals. u Atoms behave in ways to achieve noble gas configuration.

52. 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 Alkali earth metals form +2 ions.

53. 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!!!

54. 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.

55. 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.

56. 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.

57. 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.

58. 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.

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

60. 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

61. Size of Isoelectronic ions u Iso - 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

62. Size of Isoelectronic ions u Positive ions have more protons so they are smaller. Al +3 Mg +2 Na +1 Ne F -1 O -2 N -3

63. Electronegativity

64. u The tendency for an atom to attract electrons to itself when it is chemically combined with another element. u How fairly it shares electrons. u Big electronegativity means it pulls the electron toward it.

65. 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.

66. Periodic Trend u Metals are at the left end. 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. u High electronegativity.

67. Ionization energy, electronegativity Electron affinity INCREASE

68. Atomic size increases, shielding constant Ionic size increases