Chapter 7 Lecture

Chapter 7 Lecture - Chapter 7 Periodic Properties All...

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Chapter 7: Periodic Properties • All exercises as you read the text. • End of Chapter suggested exercises: – 7.2, 7.5, 7.9, 7.11, 7.13, 7.15, 7.19, 7.21, 7.23, 7.25, 7.28, 7.29, 7.31, 7.33, 7.37, 7.39, 7.43, 7.45, 7.49, 7.53, 7.57, 7.61, 7.67, 7.69, 7.71, 7.73, 7.81, 7.85, 7.86, 7.91, 7.92, 7.99, 7.108
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http://www.coolschool.ca/lor/CH11/unit5/U05L03.htm Mendeleev’s Periodic Table -Mendeleev and Meyer both came up with similar classification schemes; Moseley arranged by atomic numbers rather than atomic mass -When arranged by increasing atomic mass, elements in the same group generally have similar chemical properties. -However, properties are not identical Development of the Periodic Table
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Development of Periodic Table - Mendeleev predicted the discovery of germanium (which he called eka-silicon) as an element with an atomic weight between that of zinc and arsenic, but with chemical properties similar to those of silicon.
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Periodic Trends • In this chapter, we will rationalize observed trends in – Sizes of atoms and ions. – Ionization energy. – Electron affinity.
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Effective Nuclear Charge (Z eff ) • In a many-electron atoms, all electrons are attracted to the nucleus • Electrons are repelled by other electrons • Inner electrons screen outer electrons from the positive charge of the nucleus • The screening effect increases going down a group (column) and stays constant across a period (row) • The nuclear charge that an electron experiences depends on both factors. Z eff = Z S Z = atomic number (# of protons) S = screening constant (S is usually close to the number of inner electrons)
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Effective Nuclear Charge • Effective nuclear charge ( Z eff ): Z eff = Z S Z = atomic number S = screening constant (S is usually close to the number of inner electrons) Na atom: 1s 2 2s 2 2p 6 3s 1 = [Ne] 3s 1
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• Radial probability = probability of finding the electron at a given distance from the nucleus • However, due to radial probability, a 2s electron will more likely be closer to the positive nucleus than a 2p electron • This causes the energy ( l ) of a s orbital ( l =0) to be lower than that of a p orbital ( l =1) • Due to the 2s electron being so close to the nucleus there will be less shielding, and less shielding = greater effective nuclear charge • For orbital energies: ns < np < nd Effective Nuclear Charge http://www.physchem.ox.ac.uk/~hill/tutorials/qm1_tutorial/atomorb/index.html; http://www.tannerm.com/Quick_atom/A5.htm; and http://www.everyscience.com/Chemistry/Inorganic/Atomic_Structure/c.1101.php; http://winter.group.shef.ac.uk/orbitron/
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This note was uploaded on 04/02/2008 for the course CHEM 101 taught by Professor Farahh during the Fall '02 term at UNC.

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Chapter 7 Lecture - Chapter 7 Periodic Properties All...

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