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

2. Trends in Atomic Size u Influenced by two factors. u Energy Level Higher energy level is further away. u Charge on nucleus More charge pulls electrons in closer.

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

4. Periodic Trends u As you go across a period the radius gets smaller. u Same energy level. u More nuclear charge. u Outermost electrons are closer. NaMgAlSiPSClAr

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

6. Ionization Energy u The amount of energy required to completely remove an electron from an atom. u Removing one electron makes a +1 ion. u The energy required is called the first ionization energy.

7. Ionization Energy u The second ionization energy is the energy required to remove the second electron. u Always greater than first IE. u The third IE is the energy required to remove a third electron. u Greater than 1st of 2nd IE.

8. What determines IE u The greater the neg. charge the greater IE. u When atoms gain or lose electrons they become IONS u IE increases across a period Nucleus is holding elect. tighter u IE decreases going down a group Farther from nucleus u graphing

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

10. Shielding u The electron on the outside energy level has to look through all the other energy levels to see the nucleus. u A second electron has the same shielding.

11. Group trends u As you go down a group IE decreases because u The electron is further away. u More shielding.

12. Periodic trends u All the atoms in the same period have the same energy level. u Same shielding. u Increasing nuclear charge u So IE generally increases from left to right.

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

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

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

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

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

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

19. First Ionization energy Atomic number H He Li Be B C N O u Breaks the pattern because removing an electron gets to 1/2 filled p orbital

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

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

22. First Ionization energy Atomic number H He Li Be B C N O F Ne l Na has a lower IE than Li l Both are s 1 l Na has more shielding l Greater distance Na

23. First Ionization energy Atomic number

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

25. 2nd Ionization Energy u For elements that reach a filled or half filled orbital by removing 2 electrons 2nd IE is lower than expected. u True for s 2 u Alkali earth metals form +2 ions.

26. 3rd IE u Using the same logic s 2 p 1 atoms have an low 3rd IE. u Atoms in the aluminum family form + 3 ions. u 2nd IE and 3rd IE are always higher than 1st IE!!!

27. Electron Affinity u The energy change associated with adding an electron. u Easiest to add to group 7A. u Gets them to full energy level. u Increase from left to right atoms become smaller, with greater nuclear charge. u Decrease as we go down a group.

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

29. Ionic size u Anions form by gaining electrons. u Anions are bigger that the atom they come from. u Nonmetals form anions. u Anions of representative elements have noble gas configuration.

30. Configuration of Ions u Ions always have noble gas configuration. u Na is 1s 1 2s 2 2p 6 3s 1 u Forms a +1 ion - 1s 1 2s 2 2p 6 u Same configuration as neon. u Metals form ions with the configuration of the noble gas before them - they lose electrons.

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

32. Group trends u Adding energy level u Ions get bigger as you go down. Li +1 Na +1 K +1 Rb +1 Cs +1

33. Periodic Trends u Across the period nuclear charge increases so they get smaller. u Energy level changes between anions and cations. Li +1 Be +2 B +3 C +4 N -3 O -2 F -1

34. Electronegativity

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

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

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