Periodicity

1. Periodicity

2. Atomic Size } • Atomic Radius = half the distance between two nuclei of a diatomic molecule. Radius

3. Trends in Atomic Size • Influenced by three factors. • Energy Level • Higher energy level is further away. • Charge on nucleus • More charge pulls electrons in closer. • Shielding • Layers of electrons shield from nuclear pull.

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

5. 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.

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

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

8. Table of Atomic Radii

9. 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.

10. 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.

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

12. 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.

13. 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 of 2nd IE.

14. Symbol First Second Third 11810 14840 3569 4619 4577 5301 6045 6276 1312 2731 520 900 800 1086 1402 1314 1681 2080 H HeLi BeB C N O F Ne 5247 7297 1757 2430 2352 2857 3391 3375 3963

15. What determines IE • The greater the nuclear charge the greater IE. • Distance from nucleus increases IE • Filled and half filled orbitals have lower energy, so achieving them is easier, lower IE. • Shielding

16. Group trends • As you go down a group first IE decreases because • The electron is further away. • More shielding.

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

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

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

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

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

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

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

24. 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

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

26. 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

27. 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

28. First Ionization energy Atomic number

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

30. Electron Affinity - the energy change associated with the addition of an electron • Affinity tends to increase across a period • Affinity tends to decrease as you go down • in a period Electrons farther from the nucleus experience less nuclear attraction Some irregularities due to repulsive forces in the relatively small p orbitals

31. Table of Electron Affinities

32. Electronegativity

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

34. Group Trend • The further down a group the farther the electron is away and the more electrons an atom has. • More willing to share. • Low electronegativity.

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

36. Ionization energy, electronegativity Electron affinity INCREASE

37. Atomic size increases, shielding constant Ionic size increases

38. Another Way to Look at Ionization Energy

39. Yet Another Way to Look at Ionization Energy

41. Summary of Periodic Trends