Lecture 6

Lecture 6 -...

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Quantum Mechanics, Atomic Structure, and Chemical Bonding Chapter 12: reading sections 12.1 through 12.13 Homework in chapter 12: 21,23, 25, 31, 33, 35, 37, 41, 49, 55 Chapter 13: reading section 13.1 Chapter 14: reading Sections 14.1 through 14.3 omework in chapter 14: odd problems between 14.29 and 14.43. Homework in chapter 14: odd problems between 14.29 and 14.43.
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Summary of Electronic Configurations An orbital is defined by three quantum numbers: n, l, m l . Shell : All orbitals with same value of n, e.g. 2s, 2p x , 2p y and 2p z Subshell : All orbitals with same value of n and l, e.g. 2p x , 2p y and 2p z . Each shell contains n subshells. Each subshell contains (2l + 1) orbitals. Total number of orbitals = n 2 . Each orbital may contain a maximum of two electrons with opposite spins (m s = ±1/2). Total number of electrons in a shell = 2n 2 .
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The Aufbau (“Building Up”) Principle 3s < 3p < 3d 4s at Z = 20, Ca (life becomes complex!)
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s (1) p (3) d (5) f (7) Summary of Orbital Shapes
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The Aufbau Principle: Construction of the Periodic Table
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The Aufbau Principle: Construction of the Periodic Table Hund's rule: electrons go into a second 2p orbital with parallel spin.
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The Aufbau Principle: Construction of the Periodic Table Filling the 2p subshell produces another stable configuration of electrons which serves as the core shell of the third row: symbol [Ne]
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Chapter 13 Bonding: General Concepts We now have an overview of the elements in the Periodic Table and the composition of the compounds that they form. We now need to understand more about the bonding that is responsible for allowing the elements to combine and form compounds. There are two fundamental types of bonding:
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Lecture 6 -...

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