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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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
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.
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
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 ; ; and ;

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Effective Nuclear Charge Example What is the effective nuclear charge on a 4s valence electron of Mg?
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• Spring '02
• Farahh

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