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Big Picture
Electrons are tiny, negatively charged particles found in regions called orbitals around the nucleus of an atom. Each
element has a different number of electrons, and each electron in the atom has its own unique set of quantum
numbers. The electrons and the neutron interact and form the most stable arrangement possible. An atom’s electron
configuration is the arrangement of the electrons.
Key Terms
Chemistry
Electron Configuration
Ground State: Lowest energy state. All electrons are in the lowest energy orbitals.
Electron Configuration: Notation used to describe the way an atom’s electrons are arranged into orbitals. Electrons
are listed in order from lowest energy to highest.
Aufbau Principle: Electrons will fill the lowest energy state available within an atom and only occupy a higher energy
state if all lower states are full.
Pauli Exclusion Principle: Electrons cannot have the same four quantum numbers within the same atom.
Hund’s Rule: Each orbital in a set must have one electron before it gets a second.
Three Rules
There are three rules for determining the ground state electron configuration of an atom: the Aufbau principle, the
Pauli exclusion principle, and Hund’s rule.
1.When placing an electron into an orbital, follow the Aufbau principle (also called the building-up principle): an
electron occupies orbitals in order from lowest to highest energy.
2.According to the Pauli exclusion principle, no more than two electrons can be placed in any orbital.
3.Before a second electron can be placed in any orbital, follow Hund’s rule: all the orbitals of that sub-level must
contain at least one electron before a second electron can be added.
Atoms are not really built by adding protons and electrons one at a time! The Aufbau principle is just a way
to determine the electron configuration.
Orbital Filling Diagram
A convenient visual way to represent the arrangement of electrons in an atom is the orbital filling diagram. In an orbital
filling diagram:
• Each orbital is represented as a square (or a circle).
• Electrons are represented as arrows drawn inside the square (or circle).
• The arrow can point either up or down.
• Each sublevel is labeled by the principal energy level and sublevel.
• Example: 1s has principal energy level (n) = 1 and sublevel (l) = 0
The Aufbau principle, Pauli exclusion principle, and Hund’s rule can be used to complete an orbital filling diagram.
Aufbau Principle
The diagram to the right shows the energy levels of
various atomic orbitals. The diagram shows that:
• Orbitals
of greater energy are higher up on the
Image credit
diagram.
in the same sublevel of a principal energy
level are drawn next to each other horizontally and
are of equal energy.
• The s sublevel within a principal energy level is always
the lowest in energy.
• The
energy levels of one principal energy level can
overlap with the energy levels of another principal
energy level.
• Example:
the 4s sublevel is lower in energy than
the 3d sublevel
Image Credit: CK-12 Foundation, CC-BY-NC-SA 3.0
This guide was created by Steven Lai, Rory Runser, and Jin Yu. To learn more
about the student authors, visit http://www.ck12.org/about/about-us/team/
interns.
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v1.1.12.2012
Disclaimer: this study guide was not created to replace
your textbook and is for classroom or individual use only.
• Orbitals
Chemistry
Electron Configuration
cont .
Orbital Filling Diagram (cont.)
An easy way to remember the atomic sublevels in increasing energy is
to draw out a diagram like the one shown on right:
Follow the red arrows from top to bottom to figure out the filling order:
1s, 2s, 2p, 3s, 3p, 4s, 3d, etc.
Image Credit: Sharayanan, CC-BY-SA 2.5
Pauli Exclusion Principle
Hund’s Rule
Following the Pauli exclusion principle, each square (or
circle) in an orbital filling diagram can have at most two
arrows. The arrow pointing upward represents one spin
direction, and the arrow pointing downward represents
the other spin direction.
Hund’s rule minimizes the natural repulsive forces
between electrons. Each box (or circle) in the same
sublevel must contain one arrow before a second arrow
can be added.
It is WRONG to draw both arrows pointing up or
both arrows pointing down!
Electron Configuration Notation
An easy way to show the electron configuration of an atom is to write out for every sublevel containing an electron:
• The principal energy level (n)
• The letter designation of the sublevel (s, p, d, f)
• A number written as a superscript indicating the number of electrons in that sublevel
There are two ways to order the sublevels. One way is to list the sublevels in the same order as the Aufbau filling order.
The other way is to group sublevels within the same principal energy levels together. Check with your teacher which
convention you are expected to follow.
The number written as a superscript is not an exponent.
Examples:
Helium: Z = 2,
Fluorine: Z = 9,
Copper: Z = 29,
or
Adding up the numbers written as superscripts should equal Z, the atomic number, for a neutral atom.
Exceptions
There are exceptions to the three rules for determining the electron configuration. Some atoms with d and f orbitals
do not follow the Aufbau principle. Chromium and copper are two exceptions.
Chromium: Z = 24
• Configuration following Aufbau principle:
• Actual configuration:
or
or
Reason: The d sublevel can hold up to 10 electrons, and a half-filled sublevel is more stable than the Aufbau predicted
configuration.
Copper: Z= 29
• Configuration following Aufbau principle:
• Actual configuration:
or
or
Reason: The d sublevel is filled - a filled sublevel is more stable than the Aufbau predicted configuration.
It is not important to memorize all the exceptions to the Aufbau principle, but it is important to understand why
these exceptions happen.
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Common Problems
Writing Electron Configuration
• Write full electron configurations.
• Write noble gas configurations.
Example: Write the full electron configuration and noble gas configuration of zinc.
1.To write the full configuration, start at hydrogen on the periodic table and write down every orbital until
you reach zinc. Adding up the numbers written as superscripts should equal 30, the atomic number of zinc.
or
.
Check with your teacher which convention you are expected to follow.
2.Using the noble gas method, go back to the last noble gas before zinc. In this case, it is argon.
3.Write the noble gas abbreviation at the beginning, then write out the electron orbitals until you reach the element.
• Zinc has 2
orbitals, and 10
orbitals.
or
Drawing Electron Configuration Diagrams
Example: Draw the electron configuration diagram for chlorine.
1.Determine the electron configuration.
2.Draw circles. Each one represents an orbital, each of which can hold two
electrons. Remember that an s shell has one orbital, a p shell has three
orbitals, a d shell has five orbitals, and an f shell has seven orbitals.
Orbitals should go from left to right and written from bottom to top, with
low-energy orbitals at the bottom. Orbitals obey Aufbau principle.
3.Label each set of circles with the correct shell name.
4.Fill in the electrons. Remember Hund’s rule: you need to put one electron
in each orbital in a shell before you put another one in. Each electron is
represented by an arrow, so start out with arrows pointing up, then write
them pointing down.
Image Credit: CK-12 Foundation, CC-BY-NC-SA 3.0
Determining Quantum Numbers
Example: What are the four possible quantum numbers of gold’s (79) outermost electron?
1.Identify the principal quantum number (n). This is given by the coefficient of the final orbital. Generally, it is the
number of the row the element is in. However, if the element is in the d-block, then it is one less than the row
number. If it is in the f block, then it is two less than the row number. Gold is in the sixth row in the d-block. So in
this case, n = 5.
2.Next, identify the angular momentum quantum number (l). This is the orbital block that the electron occupies. In
this case, it occupies the d-block. For the s-block, l =0, for the p-block, l =1, for d-block, l =2, and for f-block,
l =3. Thus, l =2.
3.Identify the magnetic quantum number, ml. This designates the specific orbital within the energy block that the
electron occupies. Beginning on the leftmost element in this block, take the negative number of l, and count up
until you hit the positive value for l. Start back over at – l if needed. Thus, count -2, -1, 0, 1, 2, -2, -1, 0, 1. Here,
you hit gold, so ml =1.
4.Finally, find the magnetic spin number (ms). It is either positive or negative 1/2. If you found the magnetic quantum
number on the first count, ms = +1/2, and if you found it on the second count, it is -1/2. Since we had to count
from -2 to 2 twice, ms = -1/2.
The actual order the electrons fill the orbitals - the way values are assigned to ml - is arbitrary. Similarly, there
is no reason for the first electron in an orbital to have a positive spin number instead of a negative number.
The rules provided here help to ensure Hund’s rule and the Pauli exclusion principle are consistently followed.
Notes
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Chemistry
Electron Configuration Problem Guide