Trends the Periodic Table Trends more than 20
Trends & the Periodic Table
Trends • more than 20 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
When you’re done it will look like this so leave room for writing!
Going down column 1: 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 increasing # energy levels as go down
Increasing number of energy levels
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
Increasing Atomic Radius Increasing number of energy levels
Cs has more energy levels, so it’s bigger previous | index | next Li: Group 1 Period 2 Cs: Group 1 Period 6
As we go across, elements gain electrons, but they are getting smaller! Family IA or 1 IIA or 2 IIIA or 13 IVA or 14 VA or 15 VIA or 16 VIIA or 17 VIIIA or 18 Element Li Be B C N O F Ne Configuration 2 -1 2 -2 2 -3 2 -4 2 -5 2 -6 2 -7 2 -8
Increasing number of energy levels Increasing Atomic Radius Decreasing Atomic Radius
previous | index | next
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
previous | index | next across row size tends to decrease as go a bit because of greater PPP “proton pulling power”
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-)
previous | index | next +7 +1 Calculate “effective nuclear charge” • # protons minus # inner electrons
What the inner electrons do…. They Shield the charge felt by the valance electrons.
previous | index | next H and He: only elements whose valence electrons feel full nuclear charge (pull) NOTHING TO SHIELD THEM
Increasing number of energy levels Increasing Atomic Radius Decreasing Atomic Radius Increased Electron Shielding
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
Positive ions (cations) • Formed by loss of electrons • Cations always smaller than parent atom 2 e 8 e 8 e 2 e Ca Ca 8 e 8 e 2 e Ca+2
Negative ions or (anions) • Formed by gain of electrons • Anions always larger than parent atom
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
OR
Ionization Energy • = amount energy required to remove a valence electron from an atom in gas phase • 1 st ionization energy = energy required to remove the most loosely held valence electron (e- farthest from nucleus)
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
Increased Ionization Energy (harder to remove an electron) Increasing number of energy levels Increasing Atomic Radius Increased Electron Shielding Decreasing Atomic Radius Decreased Ionization Energy (easier to remove an electron)
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
Increasing number of energy levels Increasing Atomic Radius Increased Electron Shielding Decreased Ionization Energy (easier to remove an electron) Increased Electronegativity Increased Ionization Energy (harder to remove an electron) Decreasing Atomic Radius Decreased Electronegativity
Reactivity of Metals • judge reactivity of metals by how easily give up electrons (they’re losers)
Increasing number of energy levels Increasing Atomic Radius Increased Electron Shielding Decreased Ionization Energy (easier to remove an electron) Decreased Electronegativity Increased Ionization Energy (harder to remove an electron) Decreasing Atomic Radius More metallic Most reactive metal = Fr (the most metallic)
Reactivity of Non-metals • judge reactivity of non-metals by how easily gain electrons (they are winners)
Increased Ionization Energy (harder to remove an electron) Increasing number of energy levels Increasing Atomic Radius Decreasing Atomic Radius Increased Electron Shielding Decreased Ionization Energy (easier to remove an electron) Decreased Electronegativity Increased Electronegativity More metallic Most reactive metal = Fr (the most metallic) Most Reactive Nonmetal = F Nonreactive BACK
Allotropes • Different forms of element in same phase – different structures and properties • O 2 and O 3 - both gas phase – O 2 (oxygen) - necessary for life – O 3 (ozone) - toxic to life • Graphite, diamond: – both carbon in solid form
- Slides: 35