Chapter 8 kotz lecture

Chapter 8 kotz lecture - Atomic Electron Configurations and...

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1 Atomic Electron Configurations and Periodicity Electron Spin The 4 th quantum number is known as the “spin quantum number” and is designated by m s . It can have the value of either + ½ or – ½ It roughly translates to refer to the magnetic spin orientation of an electron in an atom (although remember that the electron exists as a wave of probability, not a spinning particle). The spinning electron has an associated magnetic field similar to the Earth’s magnetic field. A complete description of an electron in an atom must have four quantum numbers: n, l , m l , m s Depending on the arrangement of the electrons in an atom, the atom may be paramagnetic or diamagnetic . Paramagnetic atoms tend to be attracted to an external magnetic field. As we will see, these atoms have one or more “unpaired” electrons in the atom. Diamagnetic atoms are slightly repelled by an external magnetic field. There is a third type known as ferromagnetic (iron, nickel, cobalt, neodymium, and certain alloys). These substances have permanently aligned electron spins in “domains” of the substance. Circles represent domains. In a) (paramagnetic), the domains are not aligned until an external magnetic field is present. In b) ferromagnetic, the domains are aligned even in the absence of an external magnetic field. There are three ways of representing the electrons in an atom. 1. Spectroscopic notation (spdf notation) given by: 1s 2 2s 2 2p 2 , etc. The leading numbers are the n numbers, the letters are the l numbers and the superscripted number gives the total number of electrons within that suborbital. 2. Condensed spectroscopic notation given by: [Ne]3s 2 3p 4 Only the higher energy electrons are explicitly given. The core electrons are represented by the noble gas symbol in brackets. 3. Orbital box diagrams. Given by: Where the arrows denote m s
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2 Two other important rules to keep in mind besides the Pauli exclusion principle: The Aufbau (“building up”) principle: Electrons should be filled in using the lowest energy state possible. Hund’s Rule: When filling in electrons in a sublevel (such as the p sublevel), one electron should be placed in each orbital with parallel spin orientations before pairing them up. Orbital Energy and Electron Accomodation When filling orbitals of the same n+ l value, the lower n value is filled first Orbital Filling H – 1s 1 He – 1s 2 Li – 1s 2 2s 1 Be – 1s 2 2s 2 B – 1s 2 2s 2 2p 1 C – 1s 2 2s 2 2p 2 N – 1s 2 2s 2 2p 3 O – 1s 2 2s 2 2p 4 F – 1s 2 2s 2 2p 5 Ne – 1s 2 2s 2 2p 6 Cu – 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 1 Exceptions to the Aufbau principle Note Cr and Cu
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3 1s 2s 3s 2p y 3p x 2p x 2p z 3p y 3p z 4s S P 1 2 3 “Cubbies” Example: Give the proper spectroscopic, condensed spectroscopic and
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Chapter 8 kotz lecture - Atomic Electron Configurations and...

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