New topic The Periodic Table

1. New topicThe Periodic Table The how and why

2. The Modern TableA. Organization of the P.T. • 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.

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

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

5. 18 1 • The elements in these groups are called the representative elements 2 13 14 15 16 17

6. VIIIB IIA IA VIB VIIB IIIB IVB VB 1 1A 2 2A 8A 18 13 3A 14 4A 15 5A 16 6A 17 7A VIIIA VIIA VIA IVA IIIA 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

7. B. Metals

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

9. C. Transition metals • The Group B elements

10. D. Non-metals • Dull • Brittle • Nonconductors- insulators

11. Metalloids or Semimetals • Properties of both • Semiconductors

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

14. Group 1 are the alkali metals • Group 2 are the alkaline earth metals

15. Group 17 is called the Halogens • Group 18 are the noble gases

16. 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.

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

18. Periodic trends Identifying the patterns

19. 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

20. 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

21. 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

22. 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

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

24. 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 +

25. 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

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

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 • More shielding • So the atoms get bigger Li Na K Rb

30. 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

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

33. 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

34. 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

35. 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

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

37. Group trends • As you go down a group first IE decreases because of • More shielding • So outer electron less attracted

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

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

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

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

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

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

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

45. 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

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

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

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

49. Web elements First Ionization energy Atomic number

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