Trends & the Periodic Table

1. Trends & the Periodic Table

2. Trends • more than 20 properties change in predictable waybased location of elements on PT • some properties: - anyone know where we can find these numbers?! • Density • melting point/boiling point • atomic radius • ionization energy • electronegativity

3. When you’re done it will look like this so leave room for writing!

4. Period Element Configuration 1 H 1 2 Li 2-1 3 Na 2-8-1 4 K 2-8-8-1 5 Rb 2-8-18-8-1 6 Cs 2-8-18-18-8-1 7 Fr 2-8-18-32-18-8-1 Going down column 1: increasing # energy levels as go down

5. Increasing number of energy levels

6. Atomic Radius • Atomic radius: defined as ½ distance between neighboring nuclei in molecule or crystal • Affected by 1. # of energy levels 2. Proton Pulling Power

8. Increasing number of energy levels Increasing Atomic Radius

9. previous | index | next Cs has more energy levels, so it’s bigger Li: Group 1 Period 2 Cs: Group 1 Period 6

10. Family Element Configuration IA or 1 Li 2-1 IIA or 2 Be 2-2 IIIA or 13 B 2-3 IVA or 14 C 2-4 VA or 15 N 2-5 VIA or 16 O 2-6 VIIA or 17 F 2-7 VIIIA or 18 Ne 2-8 As we go across, elements gain electrons, but they are getting smaller!

11. Decreasing Atomic Radius Increasing number of energy levels Increasing Atomic Radius

12. previous | index | next

13. Why does this happen.. • As you go from left to right, you again more protons (the atomic number increases) • You have greater “proton pulling power” • Remember the nucleus is + and the electrons are - so they get pulled towards the nucleus • The more protons your have, the more Proton Pulling Power

14. previous | index | next as go across row size tends to decrease a bit because of greater PPP “proton pulling power”

15. We can “measure” the Proton Pulling Power by determining the Effectivenuclear charge • It is the charge actually felt by valence electrons • The equation Nuclear charge - # inner shell electrons (doesn’t include valance e-)

16. previous | index | next +7 +1 Calculate “effective nuclear charge” • # protons minus # inner electrons

17. What the inner electrons do…. They Shield the charge felt by the valance electrons.

18. previous | index | next H and He: only elements whose valence electrons feel full nuclear charge (pull) NOTHING TO SHIELD THEM

19. Decreasing Atomic Radius Increasing number of energy levels Increasing Atomic Radius Increased Electron Shielding

20. Look at all the shielding Francium's one valance electron has. It barely feels the proton pull from the nucleus. No wonder it will lose it’s one electron the easiest. No wonder it’s the most reactive metal

21. Positive ions (cations) • Formed by loss of electrons • Cations always smaller than parent atom 2e 8e 8e 8e 8e 2e 2e Ca Ca Ca+2

22. Negative ions or (anions) • Formed by gain of electrons • Anions always larger than parent atom

23. How do you know if an atom gains or loses electrons? • Think back to the Lewis structures of ions • Atoms form ions to get a valence of 8 (or 2 for H) • Metals tend to have 1, 2, or 3 valence electrons • It’s easier to lose them • Nonmetals tend to have 5, 6, or 7 valence electrons • It’s easier to add some • Noble gases already have 8 so they don’t form ions very easily

24. OR

25. Ionization Energy • = amount energy required to remove a valence electron from an atom in gas phase • 1st ionization energy = energy required to remove the most loosely held valence electron (e- farthest from nucleus)

26. previous | index | next • Cs valence electron lot farther away from nucleus than Li • electrostatic attraction much weaker so easier to steal electron away from Cs • THEREFORE, Li has a higher Ionization energy then Cs

27. Increased Ionization Energy (harder to remove an electron) Decreasing Atomic Radius Increased Electron Shielding Increasing number of energy levels Increasing Atomic Radius Decreased Ionization Energy (easier to remove an electron)

28. Electronegativity • ability of atom to attract electrons in bond • noble gases tend not to form bonds, so don’t have electronegativity values • Unit = Pauling • Fluorine: most electronegative element = 4.0 Paulings

29. Increased Electronegativity Increased Ionization Energy (harder to remove an electron) Decreasing Atomic Radius Increased Electron Shielding Increasing number of energy levels Decreased Ionization Energy (easier to remove an electron) Increasing Atomic Radius Decreased Electronegativity

31. Reactivity of Metals • judge reactivity of metals by how easily give upelectrons (they’re losers)

32. Increased Electronegativity Increased Ionization Energy (harder to remove an electron) Decreasing Atomic Radius More metallic Increased Electron Shielding Increasing number of energy levels Decreased Ionization Energy (easier to remove an electron) Decreased Electronegativity Increasing Atomic Radius Most reactive metal = Fr (the most metallic)

33. Reactivity of Non-metals • judge reactivity of non-metals by how easily gainelectrons (they are winners)

34. Increased Electronegativity Increased Ionization Energy (harder to remove an electron) Decreasing Atomic Radius Most Reactive Nonmetal = F More metallic Increased Electron Shielding Increasing number of energy levels Decreased Ionization Energy (easier to remove an electron) Decreased Electronegativity Increasing Atomic Radius Most reactive metal = Fr (the most metallic) Nonreactive BACK

35. Allotropes • Different forms of element in same phase • different structures and properties • O2 and O3 - both gas phase • O2 (oxygen) - necessary for life • O3(ozone) - toxic to life • Graphite, diamond: • both carbon in solid form