Section 14.2 Periodic Trends

1. Section 14.2Periodic Trends OBJECTIVES: • Interpret group trends in atomic radii, ionic radii, ionization energies, and electronegativities. • Interpret period trends in atomic radii, ionic radii, ionization energies, and electronegativities.

2. Trends in Atomic Size • First problem: Where do you start measuring from? • The electron cloud doesn’t have a definite edge. • They get around this by measuring more than 1 atom at a time.

3. Atomic Size } Radius Atomic Radius = half the distance between two nuclei of atoms • in the solid state (by X-ray diffraction) • or of a diatomic molecule.

4. Atomic Size Influenced by three factors: • Energy Level • Higher energy level is further away. • Charge on nucleus • More positive charge pulls electrons in closer. • Shielding effect • The inner electrons shield the outer electrons from the nuclear charge/attraction.

5. Shielding • The electron on the outermost energy level has to look through all the other energy levels to see the nucleus. • Second electron has same shielding, if it is in the same period • Shielding Increases down a Group, and is Constant across a Period. • Shielding Across a Group is Constant, but the EFFECTIVE NUCLEAR CHARGE INCREASES.

6. Group Trends in Atomic Size H Li • As we go down a group... • each atom has another energy level, • so the atoms get bigger. • Shielding increases as well, so the nucleus has less of a hold on e-… distance is longer. • The Increased size of the Energy Levels down a group outweighs the increased nuclear charge Na K Rb

7. Periodic Trends in Atomic Size • As you go across a period, the radius gets smaller. • Electrons are in same energy level. • More nuclear charge. • Shielding is constant … not an issue. • Outermost electrons are closer. Na Mg Al Si P S Cl Ar

8. Atomic Radius Overall

9. Rb Overall Periodic Trend for Atomic Radii K Na Xe Li Atomic Radius (nm) Kr Ar Ne H He Atomic Number 10

10. Trends in 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 to remove the first electron is called the first ionization energy.

11. Ionization Energy • The second IE 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.

12. Ionization Energy Table Where are the Group Effects? Symbol First Second Third 1312 2371 520 900 800 1086 1402 1314 1681 2080 H HeLi BeB C N O F Ne 11810 14840 3569 4619 4577 5301 6045 6276 5247 7297 1757 2430 2352 2857 3391 3375 3963 1A 2A 3A

13. Can we See the Effect of Nuclear Charge in this Table? Symbol First Second Third 11810 14840 3569 4619 4577 5301 6045 6276 1312 2371 520 900 800 1086 1402 1314 1681 2080 H HeLi BeB C NO F Ne 5247 7297 1757 2430 2352 2857 3391 3375 3963

14. What Affects the IE • The greater the nuclear charge, the greater the IE. • Larger positive nucleus has a greater attraction for the electrons, so the IE increases. • Greater distance from nucleus decreases IE • Electrons are further away from the attractive nucleus, and are easier to remove.

15. What Affects the IE • Filled and half-filled orbitals have lower energy, so the removal of an electron to achieve this ½ filled orbital requires unusually low IE. • Shielding effect • As Shielding increases, it is easier to “pluck” the outer electron, so the IE would decrease. Let’s look at Group & Period Trends for IE

16. H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87 Group trends on IE • As you go down a group, first IE decreases because... • The electron is further away. • More shielding.

17. Periodic trends on IE • Across the representative elements, the atoms are in the same period & have the same energy level. • Same shielding. • But, increasing nuclear charge holds e-’s tighter. • So IE generally increases from left to right. • Exceptions at full and 1/2 full orbitals.

18. Summarizing 1st Ionization Energy 1st Ionization Energy Increases

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

20. First Ionization Energy He • Li has lower IE than H • more shielding • further away • these outweigh greater nuclear charge H First Ionization energy Li Atomic number

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

22. First Ionization Energy He • B has lower IE than Be • same shielding • greater nuclear charge • By removing an electron we make the s orbital half-filled H First Ionization energy Be B Li Atomic number

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

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

25. First Ionization Energy He N • Breaks the pattern, because removing an electron leaves 1/2 filled p orbital O C H First Ionization energy Be B Li Atomic number

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

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

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

29. He First Ionization Energy Ne Ar Kr First Ionization energy Li Na K Atomic number

30. What’s the Driving Force? • Full Energy Levels require lots of energy to remove their electrons. • Noble Gases have full orbitals. • Atoms behave in ways to achieve noble gas configuration.

31. 2nd Ionization Energy • For elements that reach a filled or half-filled orbital by removing 2 electrons, 2nd IE is lower than expected. • True for s2 • Alkaline earth metals form 2+ ions.

32. 3rd Ionization Energy • Using the same logic s2p1atoms have a low 3rd IE. • Atoms in the aluminum family form 3+ ions. • 2nd IE and 3rd IE are always higher than 1st IE!!!

33. Trends in Electron Affinity What is Electron Affinity? It’s the energy change associated with adding an electron to a gaseous atom.

34. Trends in Electron Affinity • It’s easiest to add an electron to Group 7A. • It gets them to a full energy level, or completes the OCTET. • Increase from left to right: atoms become smaller, with greater nuclear charge. • Decrease as we go down a group.

35. Electron Affinity in the Periodic Table

36. Electron Affinity in 3-D

37. Trends in Ionic Size • Cations form by losing electrons. • Cations are smaller that the atom they come from. • Metals form cations. • Cations of representative elements have noble gas configuration.

38. Trends in Ionic Size • Anions form by gaining electrons. • Anions are bigger that the atom they come from. • Nonmetals form anions. • Anions of representative elements have noble gas configuration.

39. Revisiting Configuration of Ions • Ions always have noble gas configuration. • Na is: 1s22s22p63s1 • Forms a 1+ ion: 1s22s22p6 • Same configuration as Neon. • Metals form ions with the configuration of the noble gas before them - they lose electrons.

40. Revisiting Configuration of Ions • Non-metals form ions by gaining electrons to achieve noble gas configuration. • They end up with the configuration of the noble gas after them.

41. Group Trends Li1+ • Going down a Group, you are adding energy levels • Ions get bigger as you go down. Na1+ K1+ Rb1+ Cs1+

42. Periodic Trends • Across the period, nuclear charge increases so they get smaller. • Energy level changes between anions and cations. N3- O2- F1- B3+ Li1+ C4+ Be2+

43. Size of Isoelectronic ions • Iso- means the same • Iso electronic ions have the same # of electrons • Al3+ Mg2+ Na1+ Ne F1- O2- and N3- all have 10 electrons • all have the configuration: 1s22s22p6

44. Size of Isoelectronic ions • Positive ions that have more protons would be smaller. • Increase in size from most positive to most negative N3- O2- F1- Ne Na1+ Al3+ Mg2+

45. Electronegativity • The tendency for an atom to attract electrons to itself when it is chemically combined with another element. • How fair is the sharing? • Big electronegativity means it pulls the electron towards it. • Atoms with large negative electron affinity have larger electronegativity.

46. Electronegativity Group Trend • The further down a group, the farther the electron is away, and the more electrons an atom has. • Pull/Attraction of the positive nucleus is lessened due to increased distance and Shielding. • Electronegativity decreases. • More willing to share.

47. Electronegativity Period Trend • As you move across a Period, there are the same number of energy levels, the same shielding, however … • Pull/Attraction of the positive nucleus on other’s electrons increases as the nucleus gets larger • Electronegativity Increases.

48. Electronegativity Periodic Trend • Metals are at the left of the table. • They let their electrons go easily • Low electronegativity • At the right end are the nonmetals. • They want more electrons. • Try to take them away from others • High electronegativity.

49. Electronegativity in 3-D

50. Can We Possibly Summarize all of this Stuff??? Ionization Energy, Electronegativity, and Electron Affinity Increases