Electron+Configurations

Published By: AA and 2 (Insignificant) Others Correction: By Ethan Bennis and Nandish Pathak, AA did the other one.

An electron configuration is the way in which the electrons are distributed among the various orbitals of an atom. Fortunately the Periodic Table is designed in such a way that this process isn't terribly difficult. Each period (row) gives the outer energy level on which an elements' electrons are orbiting on. Each energy level is then divided into a subshell, labeled with the letters "s," "p," "d," or "f." Here is a diagram showing the label of each area of the Periodic Table. Image from http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch6/quantum.html

Each group (column) indicates the number of electrons in that particular energy level. For a greater understanding of s, p, d, and f, read the quantum numbers page.

The first energy level has a single subshell and can only hold up to two (2) electrons, and, when filled is designated with the symbol 1s². The second energy level contains two subshells, 2s and 2p. When filled, this energy level is designated with the symbol 2s¹2p⁶. Example Electron Configurations- Nitrogen (N): 1s²2s²2p³ Beryllium (Be): 1s²2s²
 * Applying Quantum Numbers in Electron Configurations:**

The third energy level contains an important change to this pattern. Like the previous energy levels, it begins with the 3s subshell, then proceeds to the 3p subshell. At this point however, the 3d subshell is skipped, and electrons go straight to the 4s subshell. This is because there is such a small difference in energy between the 3d and 4s subshells that electrons jump straight to the outer shell. Example Electron Configurations- Iron (Fe): 1s²2s²2p⁶3s²3p⁶4s²3d⁶ Calcium (Ca): 1s²2s²2p⁶3s²3p⁶4s² This rule continuous to apply in lower periods on the periodic table as well.

Hund's Rule states that for degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.This means that electrons will occupy the maximum number of orbitals possible before pairing with other electrons with the opposite spin. code Example: Nitrogen (N): 1s²      2s²          2p³ (||)    (||)    (|  )(|  )(|  ) code As you can see, in the 2p subshell, the 3 electrons are placed as to maximize the number of unpaired electrons.
 * Hund's Rule:**

In these cases, most of which occur after the first 40 elements, an electron is transferred from a ‘s’ orbital to a ‘d’ orbital. Transfers are possible because the energy levels are so similar between the two subshells that a transfer would be easy to accomplish. Electrons are transferred from the ‘s’ orbital to the ‘d’ orbital because a half-filled or filled ‘d’ subshell is more stable in the normal state than a half-filled ‘s’ subshell. Examples: Chromium:1s2 2s2 2p6 3s2 3p6 4s1 3d5 Copper: 1s2 2s2 2p6 3s2 3p6 4s1 3d10
 * Exceptions in electron configuration:**

In addition, as the configurations start to get longer, they can be abbreviated with a simple technique: Take the Noble Gas from the previous period, and display the electron configuration of the final period which the element is on. Example: Germanium (Ge): [Ar] 4s²3d¹⁰4p²
 * Abbreviated configuration:**

Citations: All information was gathered from our heads and from our AP Chemistry Book- __Chemistry: The Central Science.__ 9th Edition.