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Periodic Table: Patterns
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John Newlands 1864 arranged elements in octaves worked for some elements, but not all
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Dimitri Mendeleev Julius Lothar Meyer both independently created a form of the periodic table
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Dimitri Mendeleev Mendeleev usually credited because he showed how useful the table could be to predict properties
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Using the periodic table The periodic table has so much information contained in its organization that it will become your most valuable resource!!!
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Periodic Law Elements are arranged by atomic number, there is a periodic repetition in their physical and chemical properties.
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Periodic Table Horizontal rows: –Called “periods” –7 periods Show the number of the shell
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Periodic Table Vertical columns: –Called “groups” –Show the valance electrons –Elements in a group have similar physical and chemical properties
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Periodic Trends We have already learned one periodic trend- the way electrons are organized in atoms
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Some exceptions Draw the electron configuration of: Cr Cu
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Single electron The energy of a single electron reflects how tightly bound that e - is to the nucleus (more negative energy)
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Effective Charge We assume that the electrons are bound by a positive charge Z But there is an effective charge that e - “feels”
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Effective Charge The effective charge (Z eff ) represents the net effect of the attraction of the nuclear charge and the repulsion of the other e - ’s
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Effective Charge An electron closer to the nucleus would feel more of the positive charge
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Shielding The “protection” by the inner electrons so that the outer electrons do not feel the full nuclear charge
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Orbital Filling Shielding helps explain why the orbitals are filled in the order that they are The lower energy orbitals (ones closer to the nucleus) are always filled first
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Orbital Shielding Orbitals to the inside of the other orbitals do a good job of shielding the outer electrons
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Polyelectronic Model The energy required to remove an electron from an atom depends on two factors:
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Two Factors The effective nuclear charge The average distance of the electron from the nucleus
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Atomic radius Half the distance between the nuclei of two atoms. (Atoms are the same and bonded together)
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Element Characteristics Radius size decreases Radius size increases
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Element Characteristics Ionization energy –Energy required to remove an electron from an atom Easier to remove an electron from group 1 than group 8 Easier to remove electrons that are farther away from the nucleus (elements at the bottom of a group) Ionization energy increases Ionization energy generally decreases
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The general trend As we go across a period from left to right, the first ionization energy increases
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Why?? Electrons added in the same principle quantum level do not completely shield the increasing nuclear charge
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Exceptions Decreases in ionization energy moving across (i.e. N to O) is because e - in 2s provide some shielding for the 2p
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Electron Affinity The energy change associated with the addition of an electron to a gaseous atom X (g) + e - --> X - (g)
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Sign for the energy Defined as energy change when electron is added if addition is exothermic then the sign is negative
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Electron affinity What is the trend going down a group? Generally becomes more positive b/c e - added at increasing distance
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Trends... Electron affinities generally become more negative from left to right across a period, there are several exceptions
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Element Characteristics Electronegativity –It is a measure of the element's ability to attract the electrons which are in a bond Electronegativity increases Electronegativity decreases
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Element Characteristics Liquids Solids More metallic More nonmetallic Periodic Table
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Periodic Table Trends: Summary Atomic radius decreases Ionization energy and Electronegativity increase
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