1. Chapter 5

2.  From Democritus to Rutherford, models of the atom have changed due to new experiments.  As technology develops, a more complete model of the atom is developed.  Rutherford’s model identified the nucleus surrounded by electrons.  His model DID NOT explain why some things glow when heated.  His model DID NOT explain the chemical properties of elements.

3.  Niels Bohr (1885-1962) was one of Rutherford’s students. He added to the model.  Proposed that an electron is found only in specific circular paths, or orbits, around the nucleus.  Studied how energy of an atom changes when it absorbs or emits light.

4.  Energy levels are fixed “paths” or energies an electron can have when orbiting the nucleus.  Close to nucleus = less energy  Further from nucleus = more energy  e- must reside on an energy level  Moving from one level to another is possible if the right energy is lost or gained.  A quantum of energy is the amount of energy required to move an electron from one level to another.

5.  Energy gained or lost in an atom is not always the same.  Energy levels are not evenly spaced.  Higher levels are closer together. Therefore less energy needed to move levels further from nucleus.

6.  Erwin Schrodinger (1887-1961) devised and solved a mathematical equation describing the behavior of the electron in a hydrogen atom.  This model comes from his mathematical solutions.  Determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus.

7.  Based on his equations, Schrodinger was also able to explain atomic orbitals.  Orbitals explain the probability of finding an electron at various locations around the nucleus.  Often thought of as the region of space in which there is a high probability of finding an electron.

8.  Used to describe the region of space with the highest probability of finding an electron.  Energy level (n) ▪ n = 1,2,3,4,5,6,7  Sublevels correspond to an orbital of a different shape ▪ Each energy level has an equal number of sublevels ▪ Denoted by letters (s,p,d,f)  Orbital ▪ Describes highest probability ▪ Contained within sublevels ▪ Shapes describe probability

9. Principal Energy Level Number of sublevels Type of Sublevel Orbitals n = 111s1 n = 222s 2p 1313 n = 333s 3p 3d 135135 n = 444s 4p 4d 4f 13571357

10.  Each energy sublevel corresponds to an orbital of a different shape, which describes where the electron is likely to be found.  These “shapes” are based on mathematical probability experiments.  “s” orbital is spherical  “p” orbital is dumbbell shaped  “d” orbital is clover shaped or dumbbell in a donut. Click Here

11.  Three Rules for Electron Configuration  Aufbau Principle – e- occupy the orbitals of lowest energy first ▪ Orbitals of any sublevel are always equal energy  Pauli Exclusion Principle – atomic orbital may describe, at most, two e- of opposite spin  Hund’s Rule – One electron enters each orbital until each orbital has an electron. Then orbitals get partners.

12.  Use the websites linked to this page to help determine electron configuration.  https://www.caymanchem.com/app/template/chem Assistant,Tool.vm/itemid/4001;jsessionid=EEAEFB4 09423347FDE326280AABDD091 https://www.caymanchem.com/app/template/chem Assistant,Tool.vm/itemid/4001;jsessionid=EEAEFB4 09423347FDE326280AABDD091  http://www.chem1.com/acad/webtut/atomic/Orbita lPT.html http://www.chem1.com/acad/webtut/atomic/Orbita lPT.html

13.  Now that you know the rules… you may MEMORIZE and use this cheat sheet. I’ll show you how to use it.  http://www.mpcfaculty.net/mark_bishop/co mplete_electron_configuration_help.htm http://www.mpcfaculty.net/mark_bishop/co mplete_electron_configuration_help.htm  Now practice!

14.  Like everything else, there are exceptions to the rules.  Some atoms are more stable when their outer shells “break the rules.”  Half-filled sublevels are not as stable as filled sublevels, but they are more stable than other configurations.  Examples are copper and chromium