Electron
Configuration
Objectives
• Describe the relationship between orbitals and
energy levels for the electrons of an atom
• Describe how to write the electron
configuration for an atom
• Explain quantum numbers and how they affect
electron configuration of atoms
• Explain why the actual electron configurations
for some elements differ from those
predicted by the Aufbau principle
Important Vocabulary
• Atomic orbital
• Quantum numbers
• Principal quantum number
• angular momentum quantum number
• Magnetic quantum number
• Electron configuration
• Pauli exclusion principle
• Hund’s rule
• Aufbau principle
Atomic Orbitals
• An atomic orbital is best though of as a region
of space in which there is a high probability of
finding an electron
• Electrons are found in orbitals within energy
levels
• Within each level, electrons occupy orbitals
that have the lowest energy
• 4 different kinds of orbitals
• s, p, d and f
s Orbitals
• Are the simplest
• Are spherical in shape
• Have the lowest energy
• Hold only 2 electrons
p Orbitals
• Are dumbbell shaped
• They can be oriented three
different ways in space
• Has more energy than a s orbital
• Each p orbital can hold 2 electrons
for a total of 6 electrons
d & f Orbitals
• Are much more complex
• There are 5 possible d orbitals which are
clover leaf shapes
• There are 7 possible f orbitals
• f orbitals have the greatest energy
• Each orbital holds a maximum of 2
electrons
Orbitals
Relationship between Levels &
Sublevels
• 1st energy level has 1 sublevel = s
So it contains 1 orbital and holds 2 electrons
• 2nd energy level has 1 s orbital and 3 p orbitals
So it can hold 8 electrons
• 3rd energy level has s, p, & d orbitals
So it can hold 18 electrons
• 4th energy level has s, p, d, & f orbitals
So it can hold 32 electrons
Periodic Table Orbitals
Quantum Numbers
• Scientists have defined the region in which
electrons can be found by using 4 quantum
numbers
• A quantum number is a number that
specifies the properties of electrons
• The principal quantum number, n, indicates the
main energy level occupied by an electron
• Values of n are positive integers
• As n increases, the electron’s distance
from the nucleus and its energy increases
orbital
Quantum Numbers Continued
• Main energy levels can be subdivided
• The sublevels are represented by the angular
momentum quantum number, l
• This number indicates the shape or type of orbital of a
particular sublevel
• Chemists use a letter code for this quantum number
•
l = 0 = an s orbital
•
l = 1 = an p orbital
•
l = 2 = an d orbital
•
l = 3 = an f orbital
Magnetic Quantum Numbers
• Symbolized by m
• Is a subset of the angular quantum number
• It also indicates the numbers and
orientations of orbitals around the nucleus
• The value of m is in whole-number values but
depends on the value of l
• 1 s orbital, 3 p orbitals, 5 d orbitals and 7 f
orbitals
Spin Quantum Number
• Symbolized by + ½ and – ½ or ↑ and ↓
• It indicates the orientation of an electron’s
magnetic field relative to an outside magnetic
field
• A single orbital can hold a maximum of 2
electrons, which must have opposite spins
Principal Quantum
Number
Angular
Momentum
Quantum Number
Magnetic
Quantum
Numbers
Spin Quantum
Numbers
Electron Configurations
• Electrons in atoms tend to assume
arrangements with the lowest possible
energies
• An electron configuration is the written
arrangement of electrons in an atom
• It shows the lowest-energy arrangement of
the electrons for an element
• It is a shorthand notation
• 3 rules determine it
Pauli Exclusion Principle
• Was established by the German chemist
Wolfgang Pauli in 1925
• It states that each orbital can hold a
maximum of 2 electrons
• In other words, no two electrons in the same
atom can have exactly the same four
quantum numbers
• In addition, if one electron has a spin quantum
number of + ½ than the other must be – ½
The Aufbau Principle
• Also helps to write the electron configuration
for an atom
• It states that electrons fill orbitals that
have the lowest energy first
• Order for filling orbitals is as follows:
• 1s ‹ 2s ‹ 2p ‹ 3s ‹ 3p ‹ 4s ‹ 3d
Hund’s Rule
• States that orbitals of the same
n and l quantum numbers are each
occupied by 1 electron before
any pairing occurs
Electron Configuration
Example
• Let’s do magnesium (Mg)
• It has 12 electrons
Electron Configuration
1s2 2s2 2p6 3s2 or
[Ne]3s2
Electron Configuration
Practice
• Let’s try a few together
• Na
• Ar
• Ca
Electron Configuration
Electron Configuration
Practice
•Now you try!
• F
• S
Exceptional Electron
Configurations
• The elements after Vanadium (atomic #23)
do not follow the Aufbau principle
• Filled energy sublevels are more stable than
partially filled sublevels.
• Exceptions to the aufbau principle are due to
subtle electron-electron interactions in
orbitals with very similar energies.
• For Example: Cu and Cr
Exceptions to Aufbau’ Principle
INCORRECT
CORRECT