The Periodic Table

1. The Periodic Table The how and why

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

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

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

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

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

7. 8A0 1A • The elements in the A groups are called the representative elements 2A 3A 4A 5A 6A 7A

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

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

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

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

12. H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87 1s1 1s22s1 1s22s22p63s1 1s22s22p63s23p64s1 1s22s22p63s23p64s23d104p65s1 1s22s22p63s23p64s23d104p65s24d10 5p66s1 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s1

13. He 2 1s2 1s22s22p6 1s22s22p63s23p6 1s22s22p63s23p64s23d104p6 1s22s22p63s23p64s23d104p65s24d105p6 1s22s22p63s23p64s23d104p65s24d10 5p66s24f145d106p6 Ne 10 Ar 18 Kr 36 Xe 54 Rn 86

14. S- block s1 • Alkali metals all end in s1 • Alkaline earth metals all end in s2 • really have to include He but it fits better later. • He has the properties of the noble gases. s2

15. Transition Metals -d block s1 d5 s1 d10 d1 d2 d3 d5 d6 d7 d8 d10

16. The P-block p1 p2 p6 p3 p4 p5

17. f6 f13 f1 f2 f3 f4 f5 f7 f8 f10 f12 f14 f11 f9 F - block • inner transition elements

18. 1 2 3 4 5 6 7 • Each row (or period) is the energy level for s and p orbitals.

19. D orbitals fill up after previous energy level so first d is 3d even though it’s in row 4. 1 2 3 4 5 6 7 3d

20. 1 2 3 4 5 6 7 • f orbitals start filling at 4f 4f 5f

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

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

23. The Shorthand • Write the symbol of the noble gas before the element. • Then the rest of the electrons. • Aluminum - full configuration. • 1s22s22p63s23p1 • Ne is 1s22s22p6 • so Al is [Ne] 3s23p1

24. More examples • Ge = 1s22s22p63s23p64s23d104p2 • Ge = [Ar] 4s23d104p2 • Hf=1s22s22p63s23p64s23d104p65s2 4d105p66s24f145d2 • Hf=[Xe]6s24f145d2

25. The Shorthand Again Sn- 50 electrons The noble gas before it is Kr Takes care of 36 Next 5s2 Then 4d10 Finally 5p2 [ Kr ] 5s2 4d10 5p2

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

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

28. Trends in Atomic Size • Influenced by two factors. • Energy Level • Higher energy level is further away. • Charge on nucleus • More charge pulls electrons in closer.

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

30. Periodic Trends • As you go across a period the radius gets smaller. • Same energy level. • More nuclear charge. • Outermost electrons are closer. Na Mg Al Si P S Cl Ar

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

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

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

34. Symbol First Second Third 11810 14840 3569 4619 4577 5301 6045 6276 5247 7297 1757 2430 2352 2857 3391 3375 3963 1312 2731 520 900 800 1086 1402 1314 1681 2080 HHeLiBeBCNO F Ne

35. Symbol First Second Third 11810 14840 3569 4619 4577 5301 6045 6276 5247 7297 1757 2430 2352 2857 3391 3375 3963 1312 2731 520 900 800 1086 1402 1314 1681 2080 HHeLiBeBCNO F Ne

36. What determines IE • The greater the nuclear charge the greater IE. • Distance form nucleus increases IE • Filled and half filled orbitals have lower energy, so achieving them is easier, lower IE. • Shielding

37. Shielding • The electron on the outside energy level has to look through all the other energy levels to see the nucleus

38. Shielding • The electron on the outside energy level has to look through all the other energy levels to see the nucleus. • A second electron has the same shielding.

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

40. Periodic trends • All the atoms in the same period have the same energy level. • Same shielding. • Increasing nuclear charge • So IE generally increases from left to right. • Exceptions at full and 1/2 fill orbitals.

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

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

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

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

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

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

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

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

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

50. Ne He • Na has a lower IE than Li • Both are s1 • Na has more shielding • Greater distance F N O C H First Ionization energy Be B Li Na Atomic number