1. Electron Configuration
Models of the Atom
Electron Configuration

2. Development of Atomic Models
Rutherford proposed an atomic model in which the electrons move around the nucleus, like the planets move around the sun.
Rutherford’s atomic model could not explain the chemical properties of elements
Niels Bohr, a student of Rutherford’s, proposed that an electron is found only in specific circular paths, or orbits, around the nucleus.

3. Bohr’s Atomic Model

4. Energy Levels
Each possible electron orbit in Bohr’s model has a fixed energy. Energy levels are the fixed energies that an electron can have.
Energy levels are similar to rungs on a ladder.
Electrons cannot occupy the space between levels.
To move from one energy level to another an electron must gain or lose energy.
A quantum of energy is the amount of energy required to move an electron from one energy level to another energy level.

5. Energy Levels Continued
The higher the electron is on the energy ladder, the farther it is from the nucleus.
Energy levels in an atom are not equally spaced.
Higher energy levels are closer together.
When the levels are closer together it requires less energy to move from one level to the next higher.

6. Quantum Mechanical Model
Similar to the Bohr model is restricts the energy of electrons to certain values.
Does not involve an exact path the electron takes around the nucleus.
Determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus.
The electron cloud can be compared to the blur created by a propeller blade.

7. Atomic Orbitals
An atomic orbital is often thought of as a region of space in which there is a high probability of finding an electron.
The energy levels of electrons are labeled by principal quantum numbers (n). n= 1,2,3,4…
Each energy level contains sublevels
Each energy sublevel corresponds to an orbital of a different shape, which describes where the electron is likely to be found.

8. Atomic Orbitals Continued
Different atomic orbitals are denoted by letters
s orbitals are spherical, p orbitals are dumbbell-shaped
The s orbitals are spherical shaped so there is only one orientation, however the dumbbell shaped p orbitals can be oriented on the x, y, or z axis

9. Different Orbital Shapes

10. The number and kinds of atomic orbitals depend on the energy sublevels.

11. Electron Arrangement in Atoms

12. Electron Configurations
Three Rules
The Aufbau principle
The Pauli exclusion
Hund’s rule

13. Aufbau Principle Electrons occupy the orbitals of lowest energy first
The orbitals for any sublevel of a principal energy level are always of equal energy
s sublevel is always the lowest-energy sublevel
The range of energy levels within a principal energy level can overlap the energy levels of another principal level
Example: the 4s orbital is lower in energy than a 3d orbital.

14. Pauli Exclusion Principle
An atomic orbital may describe at most two electrons
To occupy the same orbital, two electrons must have opposite spin
Spin can be thought of as clockwise or counterclockwise
An orbital containing electrons is written as:

15. Hund’s Rule
When using the aufbau diagram, one electron enters each orbital until all the orbitals contain one electron with the same spin direction
Electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible
Second electrons then occupy each orbital so that their spins are paired with the first electron in the orbital
Each orbital can eventually have two electrons with paired spins

16. Periodic Patters d ( n-1) Period highest energy level f (n-2)
Group total # of valence electrons
d ( n-1)
Period highest energy level
f (n-2)

17. How many outer electrons for this group
These group 1 metals have __ electrons in outer shell
These elements have -- electrons in outer shell
How many outer electrons for this group

18. Notation
There are different ways to show the electron configuration of an atom. One involves writing the energy level and the symbol for every sublevel occupied by an electron. You indicate the number of electrons occupying that sublevel with a superscript Example: Titanium 1s2 2s2 2p6 3s2 3p6 3d2 4s2 Note: The sublevels within the same principal energy level are generally written together Another way is considered a shorthand version which uses the noble gas (last element with completely filled outer shell) and writing the remainder of the unfilled sublevel Example: Titanium [Ar]3d24s2

19. Exceptions
There are exceptions to every rule Some actual electron configurations differ from those assigned using the aufbau principle because half-filled sublevels are not as stable as filled sublevels, but they are more stable than other configurations Examples: Copper Cu 1s2 2s2 2p6 3s2 3p6 3d10 4s1 Normally the 4s would be filled completely, but in this situation the Cu is more stable with this configuration