Unit 5 The Periodic Table

1. Unit 5The Periodic Table The how and why

2. Newlands -1865 • Arranged known elements according to properties & order of increasing atomic mass • Law of Octaves – pattern of chemical & physical properties repeated every 8 elements

4. Mendeleev - 1869 • Created 1st periodic table (63 elements) • Ordered by increasing atomic mass • Predicted pattern of missing elements • Started new rows and lined up columns to organize elements with similar properties • Rearranged elements so similar properties would line up correctly

6. The Modern Table • Moseley- determined the atomic number for each known element. • Elements are still grouped by properties • Similar properties are in the same column • Ordered by increasing atomic number • Added a column of elements Mendeleev didn’t know about – noble gases

8. Periodic Law • When elements are arranged in order of increasing atomic number, elements with similar properties appear at regular intervals

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

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

11. 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

12. 8A0 1A • The elements in the A groups are called the representative elements 2A 3A 4A 5A 6A 7A

13. Transition metals • The Group B elements

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

15. Three Classes of Elements • Metals • Nonmetals • Metalloids

16. Metals

17. Metals • Ductile – drawn into wires • Malleable – hammered into sheets • All solid at room temperature (except Hg- Mercury) • Conductors of heat and electricity • Families • 1 - Alkali • 2 - Alkaline Earth • Transition (B groups)

18. Group 1A are the alkali metals • VERY reactive because one valence e- • Found as compounds in nature • Not including H!

19. Group 2A are the alkaline earth metals • Still highly reactive but not as much so • as alkali metals (2 valence e-)

20. Transition Metals • The weird ones… • May lose different #s of valence electrons depending on the element with which it reacts • Less reactive than alkali or alkaline earth metals • Good conductors of electricity & heat, ductile, malleable

21. Inner Transition Metals • 1st row = lanthanides • Shiny metals similar in reactivity to alkaline earth metals • 2nd row = actinides • Unstable nuclei – all radioactive

22. Non-metals

23. Non-metals • Most are gases, some solid, and 1 liquid (Br) • More variation than metals • Families • Halogens (Group 17 or 7A) • Noble Gases (Group 18 or 8A)

24. Group 7A is called the Halogens • Most reactive non-metals – 7 valence React frequently with alkali metals

25. Group 8A are the noble gases Low reactivity, very stable, inert

26. Metalloids or Semimetals

27. Metalloids • Border the staircase between metals and nonmetals • Properties – similar to metals and nonmetals

28. Part 2Periodic trends Identifying the patterns

29. 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

30. What we will look for • Periodic trends • How those things vary as you go across a period • Group trends • How those things vary as you go down a group • Why? • Explain why these variations exist

31. Atomic Size • Where do you start measuring? • The electron cloud doesn’t have a definite edge. • Scientists focused first on diatomic elements -- measured more than 1 atom at a time

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

33. Atomic Size - Periodic Trends • The positive nucleus pulls on electrons • Periodic trend • As you move across a period, elements have more protons • The charge on the nucleus gets bigger • The outermost electrons of each element are in the same energy level • So there is more pull on the outermost electrons as you move across

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

35. Atomic Size – Group Trends • The positive nucleus pulls on electrons • Group Trend • As you go down a group, you add energy levels • Outermost electrons not as attracted by the nucleus

36. Increasing numbers of electrons between the nucleus and the valence electrons tends to decrease the force between the nucleus & the valence electrons 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 • In the same energy level (period) shielding is the same +

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

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

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

42. Atomic size increases,

44. IONIZATION ENERGY

45. It’s all about stability • Alkali metals are more stable if they lose an electron • Example • Sodium ([Ne] 3s1) • Getting rid of the 3s1 electron makes sodium more stable and creates a sodium ion (Na1+)

46. Ionization Energy • The amount of energy required to completely remove an electron from a neutral atom. • The energy required for the 1st electron is called the first ionization energy

47. Ionization Energy • The 2nd ionization energy is the energy required to remove the second electron • Always greater than 1st IE • The 3rd IE is the energy required to remove a third electron • Greater than 1st or 2nd IE

48. 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

49. Group trends • As you go down a group first IE decreases • Valence e- farther from nucleus • More shielding

50. Periodic trends • First IE increases from left to right across a period • Increased nuclear charge from added proton • Electron shielding not an issue b/c valence are all in same energy level • Exceptions at full and 1/2 full orbitals • Lower IE b/c offer stability to atom