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Ch05 Powerpoint - CHAPTER 5 CHAPTER 5 ELECTRON...

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Unformatted text preview: CHAPTER 5 CHAPTER 5 ELECTRON CONFIGURATION ANALOGY OF THE ATOM UNITED STATES NEW JERSEY MIDDLESEX COUNTY EDISON Township YOUR HOUSE ATOM ENERGY LEVEL SUBLEVEL ORBITAL LOCATION OF ELECTRON 1.Principle Energy Levels: (n = 1,2,3….7) a.k.a 1. shells Different values of n = different energy levels & different electron energies Maximum number of e­ each level holds is found by: 2n2 EX: 1st energy level (n=1) 2n2 = 2(1)2 = 2 electrons in 1st energy level Increasing energy level Increasing energy level n = 1 2 3 4 Max # 2 8 18 32 of e­s Increasing distance from nucleus 2.Sublevels: s (spherical), p (dumbbell), d, f 2. Cloud shapes The # of energy sublevels is the same as the energy level # Energy Level (n) # of Sublevels n=1 1 Type of Sublevel 1s n=2 2 2s2p n=3 3 3s3p3d n=4 4 4s4p4d4f 3.Orbitals: 3. Each sublevel has a different # of orbitals which means a different # of electrons The # of orbitals in an energy level is found by: n2 EX: 3rd energy level (n=3) 32 = 9 orbitals This makes sense because: 3rd energy level would have 3 sublevels; s sublevel with 1 orbital, p sublevel with 3 orbitals, and d sublevel with 5 orbitals. SO……1 + 3 + 5 = 9, so the formula n2 works! Sublevel(s) Sublevel(s) # of orbitals n2 Maximum # of e ­s s 1 2n2 2 s,p 4 8 s,p,d 9 18 s,p,d,f 16 32 PRACTICE Principle Energy Level 2p 2p 5 Number of Electrons in Sublevel Type of Electron Sublevel Orbitals and Electron Capacity of the First Four Principle Energy Levels Principle energy level (n) 1 Type of sublevel # of orbitals per type 4 s 1 p 3 p 3 5 s p 1 3 d 5 f 7 Maximum # of electrons (2n2) 1 d 3 1 s 2 s # of orbitals per level (n2) 1 2 4 8 9 18 16 32 HOW TO DETERMINE HOW TO DETERMINE ELECTRON CONFIGURATIONS How electrons are arranged around the nuclei of atoms 3 RULES: Aufbau Pauli Exclusion Principle Hund’s Rule Aufbau Principle: Aufbau Principle: Electrons enter orbitals of lowest energy level 1 st Start at the beginning of each arrow, and then follow it all of the way to the end. 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p 2. Pauli Exclusion Principle: Orbitals can hold only 2 electrons Each electron in the orbital has an opposite spin 1s2 3. Hund’s Rule: One electron enters each orbital until all orbitals contain one electron with parallel spins. 1s2 2s2 2p3 1s2 2s2 2p4 Bohr Model­ Bohr Model­ Line Spectra Explained Electrons can occupy specific energy levels Excited atoms can emit light Each orbital they are in has specific energy. Atomic Emission Spectrum Atomic Emission Spectrum Spectrometer: Breaks up what we see as continuous light into individual bands of light. The individual bands of light represent the exact frequency of light be given off. This corresponds to the quantum of energy that is released when an electron goes from an excited state to the ground state. Electrons move from ground state to an excited Electrons move from ground state to an excited state and back again. The bands of light are called: The bands of light are called: ATOMIC EMISSION SPECTRUM This spectrum is unique to every element. Atomic Emission Spectrum Atomic Emission Spectrum Add energy Different elements are used to give fireworks their color! To learn more about the science of fireworks see: http:// www.pbs.org/wgbh/nova/kaboom/ Electromagnetic Spectrum Electromagnetic Spectrum How Does the Atomic Emission How Does the Atomic Emission Spectrum Correspond to Electron Configuration? 1. Electrons get excited when we introduce ENERGY (fire, electricity, light) 2. Electrons jump from ground state to an excited state 3. The amount of energy needed to do this 3. The amount of energy needed to do this is a QUANTUM of energy QUANTUM: is the amount of energy needed to move an electron from one energy level to the next. 4. Electrons go back to the ground state and give off this energy in the form of light (PHOTONS) PHOTON: a quantum of light; a discrete amount of energy that behaves as a particle. E= h c / λ E= Energy h= Planck’s constant c= speed of light λ=wavelength ...
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This note was uploaded on 11/21/2011 for the course BIO 101 taught by Professor Martin during the Fall '08 term at Rutgers.

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