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Periodic Trends.ppt
- 1. Trends & thePeriodic Table
- 2. Trends • more than20 properties change in predictable way based 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 itwill look like this so leave room for writing!
- 4. 2-8-18-32-18-8-1 Fr 7 2-8-18-18-8-1 Cs 6 2-8-18-8-1 Rb 5 2-8-8-1 K 4 2-8-1 Na 3 2-1 Li 2 1 H 1 Configuration Element Period Going down column1: increasing # energy levels as go down
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- 6. Atomic Radius • Atomicradius: defined as ½ distance between neighboring nuclei in molecule or crystal • Affected by 1. # of energy levels 2. Proton Pulling Power
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- 9. previous | index| next Li: Group 1 Period 2 Cs: Group 1 Period 6 Cs has more energy levels, so it’s bigger
- 10. 2-8 Ne VIIIA or 18 2-7 F VIIAor 17 2-6 O VIA or 16 2-5 N VA or 15 2-4 C IVA or 14 2-3 B IIIA or 13 2-2 Be IIA or 2 2-1 Li IA or 1 Configuration Element Family As we go across, elements gain electrons, but they are getting smaller!
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- 12. previous | index| next
- 13. Why does thishappen.. • 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. as go acrossrow size tends to decrease a bit because of greater PPP “proton pulling power” previous | index | next
- 15. We can “measure”the Proton Pulling Power by determining the Effective nuclear 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 Calculate “effective nuclear charge” • # protons minus # inner electrons +7 +1
- 17. What the innerelectrons 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
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- 20. Look at allthe 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. 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)
- 22. •Cs valence electronlot 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 previous | index | next
- 23. Increasing number of energy levels Increasing Atomic Radius Decreasing Atomic Radius Increased Electron Shielding DecreasedIonization Energy (easier to remove an electron) Increased Ionization Energy (harder to remove an electron)
- 24. Electronegativity • ability ofatom 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
- 25. Increasing number of energy levels Increasing Atomic Radius Decreasing Atomic Radius Increased Electron Shielding Decreased Ionization Energy (easier to remove an electron) IncreasedIonization Energy (harder to remove an electron) Decreased Electronegativity Increased Electronegativity
- 27. Reactivity of Metals •judge reactivity of metals by how easily give up electrons (they’re losers)
- 28. Increasing number of energy levels Increasing Atomic Radius Decreasing Atomic Radius Increased Electron Shielding Decreased Ionization Energy (easier to remove an electron) IncreasedIonization Energy (harder to remove an electron) Decreased Electronegativity Increased Electronegativity Most reactive metal = Fr (the most metallic) More metallic
- 29. Reactivity of Non-metals •judge reactivity of non-metals by how easily gain electrons (they are winners)
- 30. Increasing number of energy levels Increasing Atomic Radius Decreasing Atomic Radius Increased Electron Shielding Decreased Ionization Energy (easier to remove an electron) IncreasedIonization Energy (harder to remove an electron) Decreased Electronegativity Increased Electronegativity Most reactive metal = Fr (the most metallic) More metallic Most Reactive Nonmetal = F Nonreactiv e BACK
- 31. How do youknow 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
- 32. Positive ions (cations) •Formed by loss of electrons • Cations always smaller than parent atom Ca 2e 8e 8e 2e Ca+2 2e 8e 8e Ca
- 33. Negative ions or(anions) • Formed by gain of electrons • Anions always larger than parent atom
- 34. Allotropes • Different formsof 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