2. Periodic Table: Patterns

3. John Newlands 1864 arranged elements in octaves worked for some elements, but not all

4. Dimitri Mendeleev Julius Lothar Meyer both independently created a form of the periodic table

5. Dimitri Mendeleev Mendeleev usually credited because he showed how useful the table could be to predict properties

7. Using the periodic table The periodic table has so much information contained in its organization that it will become your most valuable resource!!!

8. Periodic Law Elements are arranged by atomic number, there is a periodic repetition in their physical and chemical properties.

9. Periodic Table Horizontal rows: –Called “periods” –7 periods Show the number of the shell

10. Periodic Table Vertical columns: –Called “groups” –Show the valance electrons –Elements in a group have similar physical and chemical properties

12. Periodic Trends We have already learned one periodic trend- the way electrons are organized in atoms

14. Some exceptions Draw the electron configuration of: Cr Cu

16. Single electron The energy of a single electron reflects how tightly bound that e - is to the nucleus (more negative energy)

17. Effective Charge We assume that the electrons are bound by a positive charge Z But there is an effective charge that e - “feels”

18. Effective Charge The effective charge (Z eff ) represents the net effect of the attraction of the nuclear charge and the repulsion of the other e - ’s

19. Effective Charge An electron closer to the nucleus would feel more of the positive charge

20. Shielding The “protection” by the inner electrons so that the outer electrons do not feel the full nuclear charge

21. Orbital Filling Shielding helps explain why the orbitals are filled in the order that they are The lower energy orbitals (ones closer to the nucleus) are always filled first

22. Orbital Shielding Orbitals to the inside of the other orbitals do a good job of shielding the outer electrons

23. Polyelectronic Model The energy required to remove an electron from an atom depends on two factors:

24. Two Factors The effective nuclear charge The average distance of the electron from the nucleus

25. Atomic radius Half the distance between the nuclei of two atoms. (Atoms are the same and bonded together)

26. Element Characteristics Radius size decreases Radius size increases

27. Element Characteristics Ionization energy –Energy required to remove an electron from an atom Easier to remove an electron from group 1 than group 8 Easier to remove electrons that are farther away from the nucleus (elements at the bottom of a group) Ionization energy increases Ionization energy generally decreases

28. The general trend As we go across a period from left to right, the first ionization energy increases

29. Why?? Electrons added in the same principle quantum level do not completely shield the increasing nuclear charge

30. Exceptions Decreases in ionization energy moving across (i.e. N to O) is because e - in 2s provide some shielding for the 2p

31. Electron Affinity The energy change associated with the addition of an electron to a gaseous atom X (g) + e - --> X - (g)

32. Sign for the energy Defined as energy change when electron is added if addition is exothermic then the sign is negative

33. Electron affinity What is the trend going down a group? Generally becomes more positive b/c e - added at increasing distance

34. Trends... Electron affinities generally become more negative from left to right across a period, there are several exceptions

35. Element Characteristics Electronegativity –It is a measure of the element's ability to attract the electrons which are in a bond Electronegativity increases Electronegativity decreases

36. Element Characteristics Liquids Solids More metallic More nonmetallic Periodic Table

37. Periodic Table Trends: Summary Atomic radius decreases Ionization energy and Electronegativity increase