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 • Periodic Law- When the elements are arranged by increasing atomic number, there is a periodic repetition of their chemical and physical properties. • Order is in increasing atomic number

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. 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

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

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

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

23. 1 2 3 4 5 6 7 • Each row (or period) is the energy level for s and p orbitals

24. 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

25. 1 2 3 4 5 6 7 • f orbitals start filling at 4f 4f 5f

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

27. 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

28. 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

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

30. 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

31. 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

32. Part 3Periodic trends Identifying the patterns

33. 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

34. 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

35. 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

36. 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

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

38. 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

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

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

42. 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

43. 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 or 2nd IE

44. Symbol First Second Third 1181014840 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

45. What determines IE • The greater the nuclear charge the greater IE. • Filled and half filled orbitals have lower energy, so achieving them is easier, lower IE

46. Group trends • As you go down a group first IE decreases because of • The outer electron is less attracted

47. Periodic trends • All the atoms in the same period • Have Increasing nuclear charge • So IE generally increases from left to right.

48. Ionic Size • Cations are positive ions • Cations form by losing electrons • Cations are smaller than the atom they come from • Metals form cations

49. Ionic size • Anions are negative ions • Anions form by gaining electrons • Anions are bigger than the atom they come from • Nonmetals form anions

50. Electronegativity