 Knowing the position of electrons helps us with many topics:
 Why chemical reactions occur

 Why some atoms are more stable than others
 Why some elements react with only certain atoms
We want to know how many electrons an atom has, and where they’re located.

 Bohr model only begins to explain electrons…


Electrons don’t behave like normal particles.
They also act like waves.
Because they act like waves, their exact location around the nucleus cannot be
calculated.

 Quantum mechanical model of the atom
 Electrons are contained in orbitals around atoms.
Orbitals are calculated regions of space where there is a high probablility of finding an electron in the atom.
 Each orbital can hold 2 electrons
 Orbital ≠ Orbit… think firefly

 There are 4 quantum numbers that describe the energy and approximate position of an electron
 Principal Quantum Number (energy level): describes size of orbital
 Angular Quantum Number (subshells): describe the shape of the orbital
 Magnetic Quantum Number: describes the orientation in space of an orbital (x, y, z-axis)
 Spin: clockwise or counterclockwise spin

s-orbital (sphere) p-orbital (peanut) d- orbital (double peanut) f-orbital (flower)

s
5
6
7
3
4
1
2 f (n-2)
6
7 d (n-1) p
© 1998 by Harcourt Brace & Company

 Pathways that show us the order that electrons will fill an atom.
 Start from the bottom and work your way up


When working out electron configuration, it’s important to understand orbital diagrams.
 Here are 3 rules that will help us understand and interpret orbital diagrams.

 Aufbau Principle
 Electrons fill the lowest energy orbitals available.
 You cannot skip any orbitals




 Hund’s Rule
 Within a sublevel, place one e per orbital with the same spin before pairing them.
 “Empty Seat Rule”

 Each arrow you draw represents an electron. (remember the three rules!)
 Each orbital holds 2 electrons and has different shapes




The S level has only 1 orbital= total of 2 e-
The P level has 3 orbitals = total of 6 e-
The D level has 5 orbitals = total of 10 e-
The F level has 7 orbitals = total of 14 e-

Fill in the orbital diagram for
Magnesium
*Use
Smartboard
Program for
Extra
Practice

 Abbreviated orbital diagram (Boxes or dashes)

An atom’s electron configuration…
 Describes the location of electrons within the atom


Identifies the shape of the electron clouds
- regions where the electrons are held.
 Uses numbers and letters to describe electrons’ location and which electron cloud it is part of.

The “s” tells you the electron’s cloud shape.
In this case it’s a spherically shaped cloud. It is called the s sublevel. Sublevel = Shape
1s 2
The “1” tells you how far from the nucleus the electrons can go.
In this case, its in the 1 st energy level, which is the closest level to the nucleus.
The “2” simply tells you how many electrons are in this cloud.
In this case 2 electrons are creating the cloud.
Note: Every orbital can only hold 2 electrons. They must have opposite spins .

 Orbital Diagram
O
8e -
1s 2s
 Electron Configuration
1s 2 2s 2 2p 4
2p



Shorthand Configuration
This is a shorter way to do e configurations
Steps…


1) Write the symbol of the element your doing
2) Find the noble gas (group 8) that comes

 before the element you are doing.
3) To begin, put that noble gas in [brackets]
4) Starting with that noble gas, finish the electron configuration

 Longhand Configuration
1s 2 2s 2 2p 6 3s 2 3p 4
Core Electrons Valence Electrons
 Shorthand Configuration
[Ne] 3s 2 3p 4

 Example - Germanium
1
2
3
4
5
6
7
[Ar] 4s 2 3d 10 4p 2

1
2
3
4
5
6
7

 Li N Mg C



This will help
It is a way to remember the order in which electrons fill their orbitals.
When you reach the far
Left side, you reach a
“wall” and must go back to
The right hand side.