HANDS-ON CHEMISTRY
ELECTRON CONFIGURATION
INTRODUCTION
The number and position of electrons in an atom are what determines an element’s chemical properties.
The arrangement of electrons in the atomic orbitals is called the electron configuration. Up to this point, we
have only examined the orbitals in hydrogen, an element whose atoms only have one electron, so how do
we determine where to place all of the electrons in a much heavier element like gold, mercury or uranium?
THEORY
Imagine you’re reserving airplane tickets for a class trip. The plane has two seats on each side of the aisle,
and everybody in your class wants to be able to look out the window. To minimize the amount of money
you spend and make everybody as happy as possible, you’d likely follow a set of rules like these:
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Buy the cheapest possible tickets first, before buying more expensive ones.
Make sure everybody gets their own ticket. The airline won’t let passengers share seats.
Buy as many window seat tickets as you can before buying aisle seat tickets.
Electrons fill the orbitals of the atom in a similar way, generally following three rules. Just like you want to
save money and make your class happy, electrons are placed to save energy and maximize their stability.
1. The aufbau rule: Put the electron in an open orbital whose subshell has the lowest energy level.
2. Pauli exclusion principle: No two electrons can have the same quantum numbers. This means
that one orbital will have two electrons – one with ms = ½ and the other with ms = −½.
3. Hund’s rule: Put one electron in every empty orbital in a subshell before adding an electron to a
half-full orbital. All electrons in half-full orbitals will have the same spin direction.
How do we know the energy levels of the subshells? The following diagram provides the order in which
the orbitals are filled. Start at 1, then follow each arrow until you reach a subshell with an available spot for
an electron.
7
4f
5f 7 orbitals
3d
4d
5d
6d 5 orbitals
2p
3p
4p
5p
6p
7p 3 orbitals
2s
3s
4s
5s
6s
7s
5
3
1
2
1s
8
4
6
8s 1 orbital
ACTIVITY
OBJECTIVE
Students will be able to: accurately describe the configuration of electrons in an atom after applying the
rules of electron configuration and interacting with hands-on models of the structure of the atom
MATERIALS
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Electron configurator pegboard
Hex nuts
Periodic table
PROCEDURE
1. Your group will be assigned an element. Look it up in your periodic table. How many electrons
does your element have? Count out as many hex nuts as your element has electrons.
2. Look at the pegboard. Each peg represents an orbital; the column corresponds to the energy level,
while the row corresponds to the type of orbital.
3. Following the three rules discussed in Theory, decide how to place the electrons to the orbitals. For
each electron you place, fill in the table below: list which subshell you placed the electron in and its
spin, and the rule you used to decide. Use the pegboard to keep track of the electrons you place.
Electron #
Subshell
Spin (ms)
Rule
Chemists write out the electron configuration by writing a list of the symbols for each occupied orbitals,
with the number of electrons in each orbital placed as a superscript to the right of the orbital symbol. For
example, lithium has 3 electrons, with two in the 1s orbital and one in the 2s orbital. We write this configuration as “1s2 2s1.”
4. Write the electron configuration of your element in this way.
5. Now, choose another element, with an atomic number between 19 and 36. Clear all the hex nuts
from the board, then follow the three rules to find the electron configuration of this element. Write
the configuration below.
Element:
Configuration:
HOMEWORK
1. Write the electron configuration for these elements.
a. Magnesium (Mg)
b. Argon (Ar)
c. Gallium (Ga)
2. Identify which element corresponds to these electron configurations.
a. 1s2 2s2 2p6
b. 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p5
c. 1s2 2s2 2p6 3s2 3p1