Unit+2+-+Electron+Configuration


 * __Introduction to Quantum__**

//Primary Quantum Letters:// - "**n**" represents the principle energy level - "**l**" represents the orbital energy level - "**ml**" represents the magnetic quantum # - "**ms**" represents the spin quantum #

//Primary Principles/Rules:// - Aufbau Principle - Fill the **lowest** energy levels **first** -Pauli Exclusion - Each orbital can hold **at most 2 electrons** simultaneously, due to opposing spin - Hund’s Rule - Before electrons can bring to pair up, each orbital of the sub-level (s, p, d, f) must have **at least 1 electron**

//Helpful Definitions:// - Energy Level - A region around the nucleus of an atom where an electron is likely to be moving - Quantum - A small package or unit of electromagnetic energy; the amount of energy required to move an electron from its present energy level to the next higher one - Atomic Orbital - A region in space around the nucleus of an atom where there is a **high probability** of finding an electron


 * __Electron Configurations__**
 * Check out this video on Electron Configuration, it gives a simple demonstration of a couple configurations**, http://youtube.com/watch?v=fv-YeI4hcQ4

//Orbital Diagram:// -When filling in Orbital Diagrams, you use arrows to symbolize electrons. By following the principles and rules, you will work from the bottom of the orbital diagram up the energy levels until all electrons are used up. -Go to this link to practice filling in orbital diagrams. Rememeber to **follow the orbital order and principles/rules:**
 * http://www.chempractice.com/drills/java_AO.php**

- the arrangement of electrons around the nucleus of an atom in it’s ground state - there are **4 sub-levels** - S, P, D, F - each orbital allows a **specific # of electrons**
 * ** When "n" equals... ** || ** # sub-levels ** || ** sub-level ** ||
 * 1 || 1 sub-level || 1s ||
 * 2 || 2 sub-levels || 2s 2p ||
 * 3 || 3 sub-levels || 3s 3p 3d ||
 * 4 || 4 sub-levels || 4s 4p 4d 4f ||
 * 5 || 4 sub-levels || 5s 5p 5d 5f ||
 * 6 || 3 sub-levels || 6s 6p 6d ||
 * 7 || 2 sub-levels || 7s 7p ||

//Filling the Orbitals:// - when filling the orbitals, be sure to follow the Hund's Rule, Aufbau Principle and Pauli Exclusion Principle (refer to primary principle definitions for explanantions of each.)
 * 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p ...

Write out electron configuration of oxygen; Oxygen - 8e Fill starting with **1s and 2s and continue filling orbitals** until you have used up **all** the electrons. Oxygen Configuration: 1s2, 2s2, 2p4
 * Example:**

//Shorthand://

Noble gases have complete orbitals; to write the shorthand electron configuration, we must locate the novle gas of the previous energy level - So if the electron Boron with 5 electrons in total, you could write the shorthand: -- He 2s2, 2p1 (because He has 2 electrons, and therefore it takes up the first two electrons of Boron, leaving only 3 remaining to write)

Bromine - 35e __1s2, 2s2, 2p6, 3s2, 3p6__, 4s2, 3d10, 4p5 The noble gas of the previous energy level is **Argon** The underline section in Bromine configuration represents that of Argon
 * Example:**
 * Shorthand Configuration: [Ar] 4s2, 3d10, 4p5**

Practice Questions are found on the Discussion Page for the Electron Configuration Section.
Try this sample test, complete with the answer key: http://misterguch.brinkster.net/jan2006.pdf