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Published byMiles Dawson Modified over 8 years ago
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Chemical Periodicity Trends in the periodic table
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Atomic Size
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How do you measure the size of an atom? The electron cloud doesn’t have a definite edge. Can get around this by measuring covalent atomic radius.
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Atomic Size Atomic Radius = half the distance between two nuclei of a diatomic molecule. } Radius
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Atomic size is influenced by two factors: Energy Level –more occupied levels = bigger atom Charge on nucleus –More charge pulls electrons in closer
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Group trends As we go down a group electrons are added to higher energy levels so the atoms get bigger. H Li Na K Rb
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Periodic Trends As you go across a period the radius gets smaller. Same energy level, but protons pull electrons closer to nucleus. NaMgAlSiPSClAr
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Trends in Atomic Radius
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Questions: Of the following elements, which has the largest atomic radius? Why? a) Si, Mg, S b) Al, Na, Cl c) Li, Cs Mg – same energy level, smallest nuclear charge Na – same energy level, smallest nuclear charge Cs – higher occupied energy levels
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Ionic Size
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Cations Positive ions - form by losing electrons. Metals form cations Cations of representative elements have noble gas configuration. Smaller than the atom they come from because of increased attraction by nucleus for fewer remaining electrons +
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Ionic size Anions Negative ions - form by gaining electrons. Nonmetals form anions. Anions of representative elements have noble gas configuration. Larger than the atom they come from, because nuclear attraction is less for an increased number of electrons. -
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Group trends Ions get bigger as you go down (adding energy levels) Li +1 Na +1 K +1 Rb +1 Cs +1
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Periodic Trends Across the period nuclear charge increases so both cations and anions get smaller from left to right. Li +1 Be +2 B +3 C +4 N -3 O -2 F -1
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Questions: 1. Of the following ions, which ones should have the larger radius? Why? a) Na + or Cs + b) Br - or K + 2. The Mg 2+ and Na + ions have ten electrons surrounding the nucleus. Which ion would you expect to have the smaller radius? Why? Cs + It has more occupied energy levels Br - Anions are larger than cations Mg 2+ Greater nuclear charge
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Ionization Energy
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The amount of energy required to completely remove an electron from a gaseous atom (how hard it is to pull an e - off an atom) 1 st IE = removing 1 e -, 2 nd IE=removing 2 e - Na (g) Na + + e -
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Shielding The electron on the outside energy level is shielded from the nucleus by the inner electrons
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Group trends u As you go down a group first IE decreases because the electron is further away (more shielding)
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Periodic trends All the atoms in the same period have the same energy level (same shielding). As you go from left to right, nuclear charge increases so IE generally increases.
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Questions: 1. Which element in the following sets has the lowest ionization energy and why? a)B, C, F b)K, Na, Li B – same energy level, smallest nuclear charge K – electron farther away, more shielding
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Electron Affinity
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The energy given off when an electron is added to an atom how much an atom ‘wants’ an electron F (g) + e - F - (g)
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Electron Affinity Group trends Generally decreases as we go down a group because shielding increases Periodic trends Increases from left to right as atoms become smaller with greater nuclear charge
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Questions: 1. Of the following elements, which ones should have the higher electron affinity? Why? a)Se or Te b)Calcium or Chromium Se – smaller atom Chromium – greater nuclear charge
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Electronegativity
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The tendency for an atom to attract electrons to itself when it is chemically combined (BONDED) with another element. Big electronegativity means it pulls the electron towards itself.
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Group Trends The further down a group the farther the electron is away from the nucleus and the more electrons an atom has. More willing to share = low electronegativity So as you go down a group electronegativity decreases
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Periodic Trends As we go from left to right across the table, electronegativity increases, because nuclear charge is increasing and electrons are held in more strongly Metals have low electronegativity Non-metals have high electronegativities (they win the electron tug-of-war)
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Questions: 1. Which element would you expect to have the highest electronegativity? Why? 2. Put the following elements in order of increasing electronegativity: Na, P, Cl F smallest nonmetal Na, P, Cl
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