Section 3.4
Electron Arrangements
Objectives
• Express the arrangement of electrons in
atoms using electron configurations and
Lewis valence electron dot structures
New Vocabulary
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Heisenberg uncertainty principle
Atomic orbital
Quantum numbers
Principle quantum number
Electron configuration
Ground-state electron configuration
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Aufbau principle
Pauli exclusion principle
Hund’s rule
Noble gas configuration
Valence electron
Helium, neon, argon, krypton, and xenon are all in group 8A of the periodic table. These elements are all gases at room temperature and generally
do not react with other elements. They are also similar in that after atoms of each gas are excited using electricity, they emit light. As shown by
group 8A, elements in the same group have similar properties, but why is this so? This section will introduce electron configurations and Lewis
valence electron dot structures, which will help you better understand (and eventually predict) the properties of elements in certain groups.
Jump to Periodic Table
3.4 - Electron Arrangements
Electron Configurations
Atomic Orbitals
The Heisenberg uncertainty principle states that it’s impossible to know the exact velocity and position of an electron (or any other type of particle) at the same
time. Atomic orbitals are areas around the nucleus where electrons are likely to be found. There are four types of orbitals (s, p, d, and f) each with a different shape
or shapes. The shapes and orientations of the most common orbitals (the s, p, and d orbitals) are shown in Figure 3.4-1. Quantum numbers are sets of numbers
used to describe the properties of atomic orbitals and the electrons in them. They describe the shapes, sizes, and energy levels of orbitals, as well as the direction
of electrons’ spins.
Figure 3.4-1 All s-orbitals are sphere shaped. All p-orbitals are dumbbell shaped. P-orbitals may be oriented in several different
positions in space. There are several different shapes and spatial orientations that d-orbitals can take.
The principle quantum number (n) represents an atomic orbital’s size and principle (or main) energy level. (Energy levels are
also referred to as shells.) The distance of electrons from the nucleus affects their energies. The further away electrons are
from the nucleus, the more energy they have. Electrons in larger orbitals are more likely to be further from the nucleus.
Therefore, electrons in larger orbitals have more energy. As exemplified by Figure 3.4-2, because the size of an energy level 1
orbital is smaller than an energy level 2 orbital, an electron in an energy level 1 orbital is more likely to be closer to the
nucleus and have less energy than an electron in an energy level 2 orbital.
Figure 3.4-2 The higher the
energy level, the larger the
orbital(s). Electrons in larger
orbitals have more energy
because they are more likely to be
further from the nucleus.
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Principle energy levels can be divided into sublevels. (Sublevels are also referred to as subshells.) As
shown in Figure 3.4-3, sublevels are labeled using the principle energy level they belong to and their
orbital type.
Higher principle energy levels are composed of more sublevels. For example, principle energy
level 1 contains one sublevel (the 1s sublevel), while principle energy level 2 contains two
sublevels (the 2s and 2p sublevels). Each type of sublevel contains different numbers of orbitals.
Each s sublevel contains one s-orbital. In Figure 3.4-4, each box represents an orbital. Each p
sublevel contains three p-orbitals. Each d sublevel contains five d-orbitals. Each f sublevel contains
seven f-orbitals. Each orbital can hold up to two electrons. This means that the maximum number
of electrons in each type of sublevel is twice the number of orbitals in that sublevel. In Figure
3.4-4, arrows represent electrons. Each s sublevel can hold up to two electrons. Each p sublevel can
hold up to six electrons. Each d sublevel can hold up to ten electrons. Each f sublevel can hold up
to 14 electrons.
Figure 3.4-4 An s, p, d, and f sublevel are shown (using orbital
notation) holding the maximum number of electrons possible for
that type of sublevel.
Figure 3.4-3 The 1s sublevel has a
principle quantum number of 1 and an
orbital type of s.
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3.4 - Electron Arrangements
Ground-State Electron Configurations
An electron configuration describes an atom’s electron
arrangement. Systems with lower energy are more stable. A
ground-state electron configuration is an arrangement of electrons
that gives an atom the least possible energy. In other words,
ground-state electron configurations are the most stable electron
arrangements. There are three rules that govern how electrons are
arranged in ground-state electron configurations: the Aufbau
principle, the Pauli exclusion principle, and Hund’s rule.
Electrons achieve their ground state when they occupy the closest
available orbital to the nucleus because it has the lowest possible
energy. The Aufbau principle states that electrons fill available
orbitals with the least energy first. Aufbau diagrams show the
order of orbitals from least to greatest energy. As shown in Figure
3.4-5, all orbitals in the same sublevel have the same energy.
However, within a principle energy level, each sublevel has a
different amount of energy. Also, note that within a principle
energy level, s sublevels have less energy than p sublevels, which
have less energy than d sublevels, which have less energy than f
sublevels. Finally, notice that as energy increases, the order of
orbitals gets more complex because orbitals of different principle
energy levels can overlap. For example, orbitals in the 3d sublevel
actually have slightly more energy than orbitals in the 4s sublevel.
Chapter 3 - Atomic Structure
Figure 3.4-5 The Aufbau diagram shows the order of orbitals from least to greatest energy.
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The periodic table can be used as an aid to remember the order in which orbitals should be filled. This is accomplished by dividing up the periodic table into
blocks as seen in Figure 3.4-6. Groups 1A, 2A, and helium make up the s-block. The other main-group elements (groups 3A through 8A) make up the p-block.
The transition metals make up the d-block and the inner transition metals make up the f-block.
Figure 3.4-6 The periodic table can be divided into blocks. The blocks denote the orbital types of the highest occupied
orbitals of each of the elements.
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3.4 - Electron Arrangements
Moving through the periodic table from left to right and top to
bottom, each successive element has one more electron than the
element before it (for neutral atoms). For example, a hydrogen
atom has one less electron than a helium atom and a helium
atom has one less electron than a lithium atom. Also, notice that
the period number roughly corresponds to the principle energy
level, which increases moving down the periodic table. The
highest energy electrons of a neutral atom will be found in the
sublevel indicated by the period and the block of the periodic
table containing the element. For example, the highest energy
electrons of a neutral atom of lithium will be in the 2s sublevel.
Because of these periodic trends, the periodic table can be used
to determine the order in which orbitals should be filled. This
concept is shown in Figure 3.4-7.
Using period one as a starting point, the correct order for filling
Figure 3.4-7 The periodic table can be used as an aid to help remember the order in which orbitals
are filled.
orbitals is 1s then 2s, 2p, 3s, 3p, then 4s. Notice that in the d-block,
the principle energy level is actually one less than the period number, a result of the overlapping
orbitals mentioned earlier. After 3d, electrons fill 4p orbitals then 5s, 4d, 5p and then 6s orbitals. Notice
that like the d-block, the f-block does not have principle energy levels that directly correspond to the
period number. Instead, f-block orbitals have principle energy levels that are two less than the period
number. Finally, the order continues with the 4f orbitals then the 5d, 6p, 7s, 5f, and 6d orbitals.
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The Pauli exclusion principle states that each orbital can hold a maximum of two
electrons. Furthermore, electrons must have opposite spin when located in the same
orbital. The direction of the electron’s spin is represented by the direction of the arrow.
In other words, if two electrons are in the same orbital, they must be represented with
one arrow pointing down and the other pointing up. A properly filled orbital is shown
in Figure 3.4-8.
Particles with like charges repel each other. Since all electrons are negatively charged,
they tend to keep as far apart as possible. Hund’s rule states that when filling equal
energy orbitals, electrons fill each orbital singly before filling orbitals with another
electron already in them. For example, there are three p orbitals of equal energy in the
2p sublevel. When electrons fill the 2p sublevel, they will fill each p orbital individually
before filling an orbital with another electron in it. Figure 3.4-9 shows an example of
Figure 3.4-8 According to the
Pauli exclusion principle, an
orbital can hold a maximum of
two electrons, which must have
opposite spin.
the correct versus incorrect way to denote electrons in a partially full sublevel that
contains multiple orbitals.
Figure 3.4-9 The portion of the orbital diagram on the left is correctly drawn; the portion of the orbital diagram on
the right is incorrectly drawn because it does not follow Hund’s rule.
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3.4 - Electron Arrangements
Orbital Diagrams
There are two ways to represent an atom’s electron configuration. In orbital diagrams, boxes or lines are
used to represent orbitals and upward or downward facing arrows represent electrons of opposite spin.
When drawing orbital diagrams, determine the highest energy sublevel of the atom. Then, use the
aufbau principle to draw the sublevels in the order they are filled, to the highest energy sublevel of the
atom. (Remember to use the periodic table.) Finally, fill the orbitals with the correct number of electrons
for the atom (obeying the Pauli exclusion principle and Hund’s rule). Since the number of electrons is
the same as the number of protons in neutral atoms, the atomic number equals the number of electrons.
For example, see the orbital diagram of oxygen in Figure 3.4-10.
Figure 3.4-10 This is the orbital diagram of a neutral oxygen atom.
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3.4 - Electron Arrangements
Electron Configuration Notation
Writing electron arrangements in electron configuration notation is even simpler than drawing orbital
diagrams. When using electron configuration notation, there is no need to worry about denoting the
directions that electrons spin or Hund’s rule. An electron configuration simply lists an atom’s sublevels
with the number of electrons in each sublevel. First, determine the highest energy sublevel of the atom.
Next, write the sublevels in the order they are filled to the highest energy sublevel. Finally, write the
number of electrons in each sublevel as a superscript on each sublevel. For example, the ground-state
electron configuration of iron is given in Figure 3.4-11 using electron configuration notation.
Remember that the sum of the superscripts in the electron configuration should equal the total number
of electrons in the atom. Checking the sum of the superscripts serves as a quick way to review work for
major errors.
Figure 3.4-11 This is the electron configuration of a neutral iron atom.
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3.4 - Electron Arrangements
Noble Gas Configuration Notation
For elements with large numbers of electrons, using electron configuration notation requires a lot of
writing. In order to shorten this, scientists use noble gas configuration notation. Instead of writing all of
the sublevels that contain electrons, the noble gas of the previous period is written in brackets to
represent the electron configuration up to that point. This is because the electron configuration up to
that point is identical to the electron configuration of the noble gas listed. The rest of the electron
configuration is written as usual after the noble gas. As shown in the example in Figure 3.4-12, the
noble gas electron configuration of iron begins with argon since it is the noble gas in the period before
Figure 3.4-12 This is the noble gas electron
configuration of a neutral iron atom.
iron. This is because the first part of iron’s
electron configuration is identical to argon.
After argon, the configuration is written in
the same way as the full electron
configuration.
The periodic table in Figure 3.4-13
summarizes this by showing the end of the
electron configuration of each element. This
part of the electron configuration contains
valence electrons.
Figure 3.4-13 This periodic table shows the orbital notation of the highest occupied orbitals of each element.
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3.4 - Electron Arrangements
Lewis Dot Structures
Valence electrons are the electrons found in the highest occupied energy level (or
shell). They are the outermost electrons of the electron cloud. Valence electrons
are important because they establish the chemical characteristics of elements.
They are the only electrons represented in Lewis electron dot structures, where
they are symbolized as dots. As shown in Figure 3.4-14, the number of valence
electrons is the same as the group number of representative elements (or group A
elements). For example, group 7A contains elements with seven valence electrons.
The exception to this pattern in the representative elements is helium. Though it is
in group 8A, it only has two valence electrons because its valence shell (or highest
occupied energy level) can only hold two electrons. Notice that this figure
contains only representative elements. Transition metals are not representative
elements and do not always have the same number of valence electrons as their
group number.
The number of valence electrons an element has plays a major role in the behavior
of the element. Notice how elements in the same group (column) of the periodic
Figure 3.4-14 The group number of group A elements is the same as
the number of valence electrons in those elements.
table have the same number of valence electrons. This helps explain why
elements in the same group have similar chemical properties.
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3.4 - Electron Arrangements
Below are the steps for drawing an electron dot structure for an atom.
Step 1 | Identify the symbol for the element and its number of valence electrons using the periodic table.
Reasoning: The element’s symbol will be the central part of the electron dot structure, while the number of valence electrons will dictate how many dots surround
the symbol.
Step 2 | Place the corresponding number of electron dots around the symbol. Imagine that the symbol has four sides: top, right, bottom, and left. Begin by
assigning one dot per side, moving clockwise around the symbol. Then, if there are still more dots to assign, start adding a second dot to each side until all
of the valence electrons have been accounted for.
Reasoning: The four sides represent the four orbitals in the outermost energy level of an atom of a representative element. These electrons are all negatively charged,
so they repel each other and will not pair up until there is one electron in each orbital. Dot structures should never have more than two dots on any side.
Example problem | Draw the electron dot structure for phosphorus.
Step 1 | Identify the symbol for the element and its number of valence electrons using the periodic table.
Jump to Periodic Table
The symbol for phosphorus is P. Phosphorus is in group 5A so it has 5 valence electrons.
Step 2 | Place the corresponding number of electron dots around the symbol. Imagine that the symbol has four sides: top, right, bottom, and left. Begin by
assigning one dot per side, moving clockwise around the symbol. Then, if there are still more dots to assign, start adding a second dot to each side until all
of the valence electrons have been accounted for.
Does the result make sense? | The result makes sense because phosphorus has five valence electrons. The first four dots are assigned singly to each side
and the fifth dot is assigned to a side of the symbol that already has a dot.
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Lab - Electron Arrangement
Explore the electron arrangements of atoms of several elements using the electron arrangement lab. The video overview will first walk you through the
information you need to know to complete the lab. Then, you will use the lab to build electron arrangements of atoms of several different elements and see how
different parts of electron arrangements are represented in orbital diagrams, electron configurations, and Lewis dot structures. This will give you an opportunity
to apply the information you learned in this section.
As you add orbitals and electrons to build the atoms in the electron arrangement lab, observe how the orbital diagrams, electron configurations, and Lewis dot
structures are filled in correspondingly. Notice that the lab will only allow you to add orbitals and electrons in ways that obeyed the Aufbau principle, the Pauli
exclusion principle, and Hund’s rule.
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