Chemistry Notes Section 5.2
5.2 Electron Arrangement
Arrangements like this in nature are rare because they are unstable. Unstable arrangements tend to become more
stable by losing energy. If the rock were to fall over, it would have ____________ energy, but it would be
_______________ stable. Energy and stability play an important role in how electrons are configured in an atom.
All forms of matter try to stay in their _________________ possible energy state, or lowest energy configuration.
___________________________________ - the lowest possible energy state for a given substance
A. Electron Configurations
Def: the way that electrons are ________________________ in orbitals around the nucleus of the atom
3 rules for electron arrangement or configuration
Aufbau Principle, Pauli Exclusion Principle, & Hund’s Rule
B. Aufbau Principle
Electrons occupy the levels and orbitals with the _______________________ energy first
Prefer to stay at ground state: if energy added to atom and e- move up to higher energy levels, they will quickly
emit energy to move back down to ground state level.
n =1 before n = 2 or n = 3
s orbital before p and p before d
1s, 2s, 2p, 3s, 3p,
4s, 3d, 4p…
4s is exception: Lower energy than 3d
C. Pauli Exclusion Principle
Each orbital holds a pair of electrons that spin in opposite direction
How can two negative e- occupy the same orbital? Answer: the quantum property known as spin. A spinning e- acts as
a tiny magnet, with a north pole at one end and a south pole at the other. The attraction between their opposite
magnetic poles counteracts some of the e- repulsive force. If a third e- tries to enter the orbital, its spin will be the same
as one of the e- already there, and it will be repelled.
Use up and down arrows to show spin
s orbital holds 2 electrons
p has 3 orbitals: Each hold 2 electrons
Up to 6 total e-
d has 5 orbitals: Each hold 2 electrons
Up to 10 total e-
D. Hund’s Rule
You must fill each orbital in a level with one electron each before filling any orbital with two electrons
For example: 3 electrons in 2p would be
4 electrons in 2p would be
E. Writing Electron Configurations
Example: Hydrogen
Atomic number = ________
Hydrogen Electron configuration: __________________________
________ electron
Example: Helium
Atomic number = ________
________electrons
Helium electron configuration: ________________________
How many e- in the highest occupied energy level? ______
Example: Lithium
Atomic number = _________
________ electrons
Lithium electron configuration: _____________________ How many in highest? ________
Example: Carbon
Atomic Number = _________
_______ electrons
Carbon electron configuration: ______________________ How many in highest? ________
Example: Aluminum
Atomic Number = __________
________ electrons
Aluminum electron configuration: _____________________________ How many in highest? _______
Example: Nickel
Atomic number = 28
28 electrons
Nickel electron configuration: _____________________________ How many in highest? ______
Example: Bromine
Atomic Number _________
_______ Electrons
Bromine electron configuration: ______________________________ How many in highest? ________
More Examples (using periodic table)
Helium 2e_________
_____
Beryllium 4e_____________
_____
_____
_____
_____
_____ _____ _____
_____
_____
_____ _____ _____
_____
_____
_____
_____ _____ _____
_____ _____ _____
Neon 10e________________
Magnesium 12e________________
Silicon 14e_______________________
More Examples
Carbon
Fluorine
Sodium
Phosphorus
Iron
Iodine 53e1s22s22p63s23p64s23d104p65s24d105p5
1s22s22p63s23p63d104s24p64d105s25p5
Radon 86 e1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p6
1s22s22p63s23p63d104s24p64d104f145s25p65d106s26p6
F. Exceptions
A few elements are exceptions to the Aufbau Principle, because half-filled sublevels are sometimes
more stable than others, except full sublevels.
Example: Chromium, atomic number = 24
1s2 2s2 2p6 3s2 3p6 3d5 4s1 (more stable than 3d44s2)
Example: Copper, atomic number = 29
1s2 2s2 2p6 3s2 3p6 3d10 4s1 (more stable than 3d94s2)
5.2 Section Assessment
1. What are the three rules for writing the electron configuration of elements?
2. Explain why the actual electron configurations for some elements differ from those assigned using the aufbau
principle.
3. Arrange the following sublevels in order of decreasing energy:
2p, 4s, 3s, 3d, and 3p
4. Why does one electron in a potassium atom go into the fourth energy level instead of squeezing into the
third energy level along with the 8 already there?
5. Give the symbol and name of the elements that correspond to these configurations of an atom:
a. 1s22s22p63s1
b. 1s22s22p3
c. 1s22s22p63s22p2
d. 1s22s22p4
e. 1s22s22p63s23p64s1
f. 1s22s22p63s23p63d24s2
6. An atom of an element has two electrons in the first energy level and five electrons in the second energy
level. Write the electron configuration for this atom and name the element. How many unpaired electrons
does an atom of this element have?