Trends & the Periodic Table

1. Trends & the Periodic Table

2. Trends • see properties change in predictable waysbased location of elements on PT • some properties can be predicted: density melting point/boiling point * atomic radius * ionization energy • electronegativity TABLE S • anyone know where we can find these numbers

3. Periodic properties: Graph shows a repetitive pattern (Note:Doesn’t have to be a straight line)

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

5. 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 - makes sense that atoms get larger in size

6. Increasing number of energy levels

7. Atomic Radius • atomic radius: defined as ½ distance between neighboring nuclei in molecule or crystal • affected by 1. # energy levels 2. Proton Pulling Power (PPP)

8. TRENDS: atoms get larger as go down column: ↑# principal energy levels atoms get smaller as move across series: ↑PPP “proton pulling power”

9. Cs has more energy levels, so it’s bigger Li: group 1 period 2 Cs: group 1 period 6

10. Increasing number of energy levels Increasing Atomic Radius

11. 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! What is happening?

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

13. Why does this happen.. • as go from left to right, you gain more protons (atomic number increases) • results in greater “proton pulling power” • remember: nucleus is (+) and electrons are (-) so e- get pulled towards the nucleus • more protons you have, the stronger PPP

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 PPP by determining the effectivenuclear charge • this is charge actually felt by valence electrons • equation to calculate effective nuclear charge: 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. previous | index | next H and He: only elements whose valence electrons feel full nuclear charge (pull) NOTHING TO SHIELD THEM

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

19. 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

20. Ionization Energy • definition: amount energy required to remove farthest valence e- from atom • 1st ionization energy: energy required to remove most loosely held valence electron (valence e- farthest from nucleus)

21. Trends in Ionization Energy • What do you think happens to the ionization energy as go down column of PT? • As go across row? decreases increases

22. Electronegativity • ability of atom to attract electrons to itself so can form bonds with other elements (to create cmpds) • noble gases tend not to form bonds, so don’t have electronegativity values • Fluorine: most electronegative element = 4.0 Paulings • Francium: least electronegative element = 0.7 Paulings

23. Decreasing Atomic Radius Increasing number of energy levels Increasing Atomic Radius Increasing electron shielding Increasing Ionization Energy Increasing Electronegativity due to  PPP

24. previous | index | next elements in same group: farther away valence electrons are from nucleus the easier to remove them easier for Cs (top of column) to lose electrons than Li (bottom of column) so Cs is a more reactive metal!

25. previous | index | next elements in same row: easier to take away valence electrons when have less protons Li has less “proton pulling power” so easier to remove its valence electrons

26. Reactivity of Metals • metals are losers! • judge reactivity of metals by how easily give upelectronsto form (+) ions • most active metals: Fr (then Cs) • for metals, reactivityincreasesas ionization energy goes down

27. Trends for Reactivity (Metallic Character) of Metals • increases as go down column • easier to lose electrons! • decreases as go across row • more difficult to lose electrons!

28. Reactivity of Non-metals • non-metals are winners! • judge reactivity of non-metals by how easily gainelectrons • F: most active non-metal • for non-metals: • reactivity ↑ as electronegativity ↑

29. Trend for Reactivity of Non-metals:depends on PPP • increases as go across row • decreases as go downcolumn • (shielded by more inner-shell electrons)

30. 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 these than gain extra needed • non-metals tend to have 5, 6, or 7 valence electrons • it’s easier to add extra needed than to lose these • noble gases already have 8 so they don’t form ions very easily

31. 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

32. negative ions or (anions) • formed by gain of electrons • anions always larger than parent atom

33. Allotropes • different structural forms of element in same phase • different structures and properties • examples: C and O

34. Graphite and Diamond: both carbon in solid form

35. O2 (g) and O3 (g) O2 (oxygen) - necessary for life O3 (ozone) - toxic to life