Chapter 6 The Periodic Table

1. Chapter 6The Periodic Table The how and why

2. History • 1829 German J. W. Dobereiner Grouped elements into triads • Three elements with similar properties • Properties followed a pattern • The same element was in the middle of all trends • Not all elements had triads

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

4. 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

5. 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.

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

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

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

9. VIIIB IA IIA VIB VIIB IIIB IVB VB 1 1A 2 2A 8A 18 13 3A 14 4A 15 5A 16 6A 7A 17 VIIIA VIA VIIA IIIA IVA VA IB IIB 3 3B 4B 4 5 5B 6B 6 7 7B 8 8B 9 8B 10 8B 1B 11 2B 12 Other Systems

10. Metals

11. Metals • Luster – shiny. • Ductile – drawn into wires. • Malleable – hammered into sheets. • Conductors of heat and electricity.

12. Transition metals • The Group B elements

13. Non-metals • Dull • Brittle • Nonconductors- insulators

14. Metalloids or Semimetals • Properties of both • Semiconductors

15. These are called the inner transition elements and they belong here

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

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

19. 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 • So.. the properties of atoms repeat.

20. H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87 1s1 1s22s1 1s22s22p63s1 1s22s22p63s23p64s1 1s22s22p63s23p63d104s24p65s1 1s22s22p63s23p63d104s24p64d105s2 5p66s1 1s22s22p63s23p63d104s24p64d104f145s25p65d106s26p67s1

21. He 2 1s2 1s22s22p6 1s22s22p63s23p6 1s22s22p63s23p63d104s24p6 1s22s22p63s23p63d104s24p64d105s25p6 1s22s22p63s23p63d104s24p64d105s24f14 5p65d106s26p6 Ne 10 Ar 18 Kr 36 Xe 54 Rn 86

22. 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

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

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

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

26. 1 2 3 4 5 6 7 • Each row (or period) is the energy level for s and p orbitals

27. 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

28. 1 2 3 4 5 6 7 • f orbitals start filling at 4f 4f 5f

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

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

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

32. More examples • Ge = 1s22s22p63s23p63d104s24p2 • Ge = [Ar] 4s23d104p2 • Ge = [Ar] 3d104s24p2 • Hf=1s22s22p63s23p64s23d104p64f14 4d105s25p65d26s2 • Hf=[Xe]6s24f145d2 • Hf=[Xe]4f145d26s2

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

34. Practice • Write the shorthand configuration for • S • Mn • Mo • W

35. Electron configurations and groups • Representative elements have s and p orbitals as last filled • Group number = number of electrons in highest energy level • Transition metals- d orbitals • Inner transition- f orbitals • Noble gases s and p orbitals full

36. Part 3Periodic trends Identifying the patterns

37. What we will investigate • Atomic size • how big the atoms are • Ionization energy • How much energy to remove an electron • Electronegativity • The attraction for the electron in a compound • Ionic size • How big ions are

38. What we will look for • Periodic trends- • How those 4 things vary as you go across a period • Group trends • How those 4 things vary as you go down a group • Why? • Explain why they vary

39. The why first • The positive nucleus pulls on electrons • Periodic trends – as you go across a period • The charge on the nucleus gets bigger • The outermost electrons are in the same energy level • So the outermost electrons are pulled stronger

40. The why first • The positive nucleus pulls on electrons • Group Trends • As you go down a group • You add energy levels • Outermost electrons not as attracted by the nucleus

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

42. 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 • In the same energy level (period) shielding is the same +

43. Shielding • As the energy levels changes the shielding changes • Lower down the group • More energy levels • More shielding • Outer electron less attracted + Three shields No shielding One shield Two shields

44. 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

45. Atomic Size } Radius • Atomic Radius = half the distance between two nuclei of molecule

46. 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

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

48. Periodic Trends • As you go across a period the radius gets smaller. • Same shielding and energy level • More nuclear charge • Pulls outermost electrons closer Na Mg Al Si P S Cl Ar

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