THE QUANTUM THEORY AND ITS

RELATION TO ELECTRONIC
CONFIGURATION
for General Chemistry 1/ Grade 12
Quarter 2 / Week 1

NegOr_Q2_GenChem1-12_SLK Week1_v2 1 NegOr_Q2_GenChem1-12_SLK Week1_v2

 FOREWORD

This self–learning kit will serve as a guide on how to use

quantum numbers to describe an electron in an atom. It will
be your aid as you learn to describe completely the
movement and trajectories of each electron within an atom.

This self-learning kit will also serve as a guide on how to

determine the magnetic property of the atom based on its
configuration and how to draw orbital diagrams to represent
the electronic configuration of atoms. It will be your aid as
you learn electron configuration and orbital diagram based
on your periodic table.

2 NegOr_Q2_GenChem1-12_SLK Week1_v2
OBJECTIVES
At the end of the lesson, you should be able to:
K : know the concept of valence electron and quantum numbers.
S : visualized the placement of electrons by electronic
configuration.
A : Appreciate the use of quantum numbers as the fundamental
concept of the importance of electrons specifically on
bonding.

LEARNING COMPETENCIES
- Use quantum numbers to describe an e l e c t r o n in an atom
(STEM_GC11ESIIa-b-54)
- Determine the magnetic property of the atom based on its
electronic configuration (STEM_GC11ESIIa-b-57)
- Draw an orbital diagram to represent the electronic
configuration of an atom (STEM_GC11ESIIa-b-58)

I. WHAT HAPPENED
PRE-ACTIVITY/PRE-TEST

I. Identify the correct answer found in the box below and write your
answer before the number.
- Energy level - Pauli’s exclusion principle - Spin quantum number
- Neils Bohr - Valence electrons - Aufbau principle
- d orbitals - Werner Heisenberg - Sub-shells
-1 - Erwin Schrodinger - f orbitals
-0 - Uncertainty principle - Hund’s rule

_______ 1. It describes the energy of the electron to sustain its path

around the nucleus.
_______ 2. This states that no two electrons in the same atom can
have the same set of four quantum numbers.
_______ 3. It describes the spin of the electrons.
_______ 4. He describes that each energy level can only
accommodate a certain number of electrons.
_______ 5. Electrons are located at the outermost energy level.

3 NegOr_Q2_GenChem1-12_SLK Week1_v2
_______ 6. This describes how the electrons are distributed among the
orbitals.
_______ 7. It can accommodate a maximum of 10 electrons.
_______ 8. He states that the position and momentum of electrons
can be determined simultaneously.
_______ 9. These are the s, p, d, and f.
_______ 10. The lowest possible energy level.

II. WHAT I NEED TO KNOW

DISCUSSION:

Quantum Numbers
Each electron in an atom is unique. This was described by
Pauli’s exclusion principle that no two electrons can have the same
sets of quantum numbers. The quantum numbers describe the
electron in an atom. An analogy to this is similar to when you watch
a movie in a movie house. You buy a ticket, in the ticket, it indicates
the row and the seat. This is similar to the electron in an atom. The
electron has its unique location and the quantum number describes
where the electron is.

Principal Quantum Numbers

The principal quantum number (n) indicates the size of the
orbital (see Figure 1). The bigger the n is, the greater the average
distance of an electron in the orbital from the nucleus, and thus, the
larger the orbital and the atom. The principal quantum number also
indicates the main energy level occupied by an electron and takes
on positive integers (1, 2, 3, and so on) as values. For example, an
electron with n = 1 occupies the first level closest to the nucleus. More
than one electron can occupy the same energy level in the same
shell. A shell is composed of a set of orbitals that have the same
principal quantum n.

4 NegOr_Q2_GenChem1-12_SLK Week1_v2
 Figure 1. The energy level (n) is the energy of the electron to sustain its path. Note, that the lowest
energy level is 1. As n increases, the distance from the nucleus increases.

Source: https://ecampusontario.pressbooks.pub/app/uploads/sites/869/2021/04/image1-2.jpg

Azimuthal Quantum Numbers

The azimuthal or orbital or angular quantum number (ℓ),
also known as angular momentum quantum number,
corresponds to the shape of the orbital. Allowed values for ℓ
include 0 and all positive integers less than or equal to n-1. For
example, an n =1 denotes a specific orbital shape corresponding
to ℓ = 0. An n = 2 may have one of two orbital shapes
corresponding to ℓ = 0 and ℓ = 1. Letters are assigned to the
orbitals depending on the l values. Therefore, an ℓ = 0
corresponds to an s orbital; ℓ = 1 means a p orbital; and so on.
One or more orbitals with the same values of n and l are known
as subshells. For example, the shell with n = 2 has two subshells (ℓ =
0 and ℓ = 1). These subshells are referred to as the 2s and 2p
subshells. The number “2” refers to n, while s and p refer to the two
values of l. Each subshell can accommodate only a certain
number of electrons as indicated in Table 1.
Table 1. Azimuthal or orbital or angular quantum numbers describe the shape of
the electron probability in an atom. The different shapes are also known as
subshells and each has its own shape.

Source: https://study.com/academy/lesson/angular-momentum-quantum-number-definition-example-quiz.html

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Magnetic Quantum Number
The magnetic quantum number (mℓ) indicates the orientation
of an orbital around the nucleus (see Figure 2). For a particular value
of l, there will be (2ℓ + 1) possible values of ml. Hence the values of ml
are integers from –l to +l, including 0. For example, if ℓ = 0, only one
value for ml is possible; that is ml = 0. If l = 1, there are (2ℓ + 1 = 3)
possible values of mℓ which are -1, 0, and +1. The number of ml values
also gives an idea about the number of orientations of orbitals
belonging to a particular subshell. For instance, for a p orbital with ℓ =
1, three possible ml values (-1, 0, +1) imply that there are three possible
orientations of the p orbital around the nucleus.

Figure 2. The magnetic quantum number describes the orientation in space of the electrons. Notice that, s has only
1 orientation since it’s spherical. For p, have 3 orientations, d have 5 orientations, and f have 7 orientations. Each
orientation can accommodate a maximum of 2 electrons. So, s can have a maximum of 2 electrons, p can have a
maximum of 6 electrons, d can have a maximum of 10 electrons, and f can have a maximum of 10 electrons.
Source: https://byjus.com/chemistry/quantum-numbers/

Spin Quantum Numbers

The spin quantum number (ms) indicates the spins of the
electrons and may have only two possible values, +1/2 and -1/2. The
(+) and (-) signs only refer to the orientation of the spin, and not to
the electric charge (see Figure 3).

Figure 3. The spin quantum number describes the spin of the electrons. This spin creates a partial
charge of ± ½. This explains why there are 2 electrons occupying an orbital excise even though both
electrons are negatively charged.
Source: https://study.com/academy/lesson/spin-quantum-number-definition-example- quiz.html
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Rules Governing the Combination of Quantum Numbers
1. The quantum numbers n, ℓ, mℓ, are integers.
2. The principal quantum number, n, cannot be zero (0). Its lowest
value is 1. Example, n = 1, 2, 3, 4, 5, and so on
3. The azimuthal quantum number, ℓ, can have a value from 0 to n-
1
4. The highest value it can have depends on n. Example, if n = 6, ℓ
can be 0, 1, 2, 3, 4 or 5.
5. The magnetic quantum number, ml, can be any integer from –l to
+l. Example, if ℓ = 1, mℓ can be -1, 0 or +1. If ℓ = 2, mℓ can be -2, -1,
0, +1 or +2.
6. The spin quantum number, ms, can only be +1/2 or -1/2. No other
values are allowed.

Electronic configuration
Electron distribution is important in understanding chemical
bonding and chemical reactions. The step-by-step process of
arranging electrons in an atom is called electronic configuration.
Based on the quantum number, each electron is unique and has its
own quantum number. This is stated by Pauli’s exclusion principle
which states that no two electrons can have the same sets of
quantum numbers. In building an atom by adding electrons, it should
start from the lowest possible energy level which is n = 1. An analogy
to this is when building a house, it should start from the foundation
and up, it should not be possible to start the house with a roof. As seen
in the diagram, as energy level (n) increases, the number of sub-shells
that needs to be filled up also increases as well as the number of
electrons. In addition, the distance from the nucleus also increases as
the energy level increases (see figure 4).

7 NegOr_Q2_GenChem1-12_SLK Week1_v2
 - In filling electron, the green line should
be followed starting from 1s.
- Notice, that when 3p is reach, the next
to be filled up is not 3d but 4s. This is also
be seen on a higher energy level of
different orbital. This is due to the
overlapping of orbitals. And this is
experimentally been proven. Note,
that s can occupy a maximum of 2
electrons, p can occupy a maximum
of 6 electrons, d can occupy a
maximum of 10 electrons, and f can
occupy a maximum of 14 electrons.
Figure 4. Aufbau Diagram – a
mnemonic used in electronic
configuration

There are two notations used in writing the electronic

configuration, the first is spdf notation and orbital notation. Spdf
notation highlighted the number of electrons occupying a certain
energy level. On the other hand, orbital notation highlighted on the
number of electrons occupying a certain subshell or orbital. Below
are some examples of the notations used for writing the electronic
configuration.
A. Spdf notation of Helium, He (2 electrons);
Number of electron/s
2
Energy level (n)
1s Azimuthal (ℓ)
B. Orbital notation of Neon, Ne (10 electrons)

or

Note that in some sources, they are using a line instead of a box; 1 line for s subshell, 3
lines for p subshell, 5 lines for d subshells, and 7 lines for f subshells. The boxes and lines in
orbital notation represent degeneracy. Degeneracy means they are on the same
energy but differ in spatial orientation. For example, p subshell contains 3 possible spatial
orientations px, py, and pz. t the 3 p’s. The basic electron distribution is made when an
atom is in its ground state. Ground state means that an atom’s electrons occupy the
lowest possible energy levels. Representing electron distribution follows three general
rules – the Aufbau principle, the Pauli exclusion principle, and Hund’s rule.

8 NegOr_Q2_GenChem1-12_SLK Week1_v2
 The Aufbau principle states that electrons should occupy first the
orbitals with lower energy before those with higher energy. From figure
1, 1s orbital should be filled first before the 2s orbital. Take note also
that the 4s orbitals have lower energy than 3d orbital. The Pauli
exclusion principle states that no Assuming 2S orbital with 2 electrons. The
two electrons in an atom can quantum numbers for arrow up: n = 0, ℓ = 0,
possess the same set of quantum mℓ = 0, ms = +1/2. For arrow down; n = 0, ℓ =
0, mℓ = 0, ms = -1/2
numbers. This principle
emphasizes the significance of
the spin quantum numbers. If two
electrons in an atom have the
same n, l, ml they should still have Figure 2. The arrow up and arrow down
different ms values; meaning, one represents electrons and each have its own
unique quantum numbers.
electron must have ms = +1/2 and
the other must-have ms = -1/2. This principle is best explained using
orbital diagrams. Consider the two electrons in the 1s orbital. Recall
that the s orbital can accommodate up to two electrons only. Take
note of the set of quantum numbers of each electron.
Lastly, Hund’s rule of maximum multiplicity or simply Hund’s rule
suggests that the most stable arrangement of electrons in subshells is
the one with the greatest number of parallel spins. This means that
each orbital in a subshell is singly occupied before the pairing of
electrons occurs. Below is an example of applying Hund’s rule;

CORRECT:

INCORRECT:

(a) (b) (c)

Magnetic property based on electronic configuration

Magnetism is a property of materials that respond to an applied

magnetic field. Permanent magnets have persistent magnetic fields
caused by ferromagnetism, the strongest and most familiar type of
magnetism. However, all materials are influenced differently by the

9 NegOr_Q2_GenChem1-12_SLK Week1_v2
presence of a magnetic field. Some are attracted to a magnetic field
(paramagnetic); others are repulsed by it (diamagnetic); still, others
have a much more complex relationship with an applied magnetic
field (e.g., spin-glass behavior and antiferromagnetism). Substances
that are negligibly affected by magnetic fields are considered non-
magnetic, these are copper, aluminum, gases, and plastic. Pure
oxygen exhibits magnetic properties when cooled to a liquid state.

The magnetic properties of a given element depend on the

electron configuration of that element, which will change when the
element loses or gains an electron to form an ion. If the ionization of
an element yields an ion with unpaired electrons, these electrons may
align the sign of their spins in the presence of a magnetic field, making
the material paramagnetic. If the spins tend to align spontaneously in
the absence of a magnetic field, the resulting species is termed
ferromagnetic.

Determination if paramagnetic or diamagnetic

The magnetic properties of a substance can be determined by

examining its electron configuration: If it has unpaired electrons, then
the substance is paramagnetic and if all electrons are paired, the
substance is then diamagnetic. This process can be broken into three
steps:

1. Write down the electron configuration

2. Draw the orbital notation
3. Identify if unpaired electrons exist at the outermost energy level.
4. Determine whether the substance is paramagnetic or diamagnetic.
Paramagnetic if it contains at least 1 unpaired electron and
diamagnetic if it has paired electrons.

Example 1: Determine the electronic configuration of Cl using a.) spdf

notation, and b.) orbital notation.
Solution.
For a.) Find the number of electrons. For Cl atoms, there are 17 electrons

10 NegOr_Q2_GenChem1-12_SLK Week1_v2
 - Writing the spdf notation (Note, use Aufbau diagram):
1s22s22p63s23p5
For b.) Writing the orbital notation:

1S 2S 2p 2s 3p
Based on the orbital notation, the Cl contains an unpaired electron in
the 3p orbital, thus it is expected to be paramagnetic, albeit weak.

III. WHAT I HAVE LEARNED

EVALUATION/POST TEST:

I. Give the correct answers to the questions given below.

1. Quantum Numbers are solutions of _____________

a) Heisenberg’s Uncertainty Principle
b) Einstein’s mass-energy relation
c) Schrodinger’s Wave Equation
d) Hamiltonian Operator

2. Which of the following quantum number gives the shape of the

atomic orbital of the sub-shell?
a) n b) ℓ c) mℓ d) s

3. Which of the following can be the quantum numbers for an

orbital?
a) n = 4, ℓ = 4, mℓ = 3 c) n = 3, ℓ = 2, mℓ = -1
b) n = 2, ℓ = 3, mℓ = 1 d) n = 3, ℓ = 0, mℓ = -3

4. “No two electrons in an atom can have the same set of n, ℓ, mℓ,
and ms quantum numbers." This is a statement:
a) Pauli exclusion principle c) Hund’s rule
b) Aufbau principle d) Bohr’s theory

5. Element Z has the ground state electronic configuration 1s22s22p3.

Contains how many unpaired electrons in the outermost shell?
a.) 1 b) 2 c) 3 d) 4
11 NegOr_Q2_GenChem1-12_SLK Week1_v2
II. Give the ground state electron configuration for:
Element Using Spdf notation
a. Boron (5 e-)
b. Silicon (14 e-)
c. Hydrogen (1 e-)
d. Beryllium (4 e-)
e. Chlorine (17 e-)

III. Write the orbital notation by filling up the box and identify if
paramagnetic or diamagnetic
1. Sodium (11 electrons):
__ __ __ __ __ __ __ __ __
1s 2s 2p 3s 3p
2. Neon (10 electrons):
__ __ __ __ __ __ __ __ __
1s 2s 2p 3s 3p
2. Titanium (22 electrons):
__ __ __ __ __ __ __ __ __ __ __ __ __ __ __
1s 2s 2p 3s 3p 4s 3d

REFERENCES

Ebbing, Darrell, and Steven D. Gammon. General chemistry.

Cengage Learning, 2016.
"Quantum Number." Encyclopædia Britannica. Accessed October
10, 2021. https://www.britannica.com/science/quantum-
number.
BYJUS. 2021. Electron Configuration - Detailed Explanation with
Examples. [online] Available at:
<https://byjus.com/chemistry/electron-configuration/>
[Accessed 10 October 2021].
Zumdahl, Steven S. & Zumdahl, Susan A. (2007). Chemistry 7th ed.
Houghton MifflinCompany, USA. Pp106-111.
12 NegOr_Q2_GenChem1-12_SLK Week1_v2
 DEPARTMENT OF EDUCATION
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Writer
_________________________________

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DISCLAIMER

The information, activities and assessments used in this material are designed to provide accessible learning modality
to the teachers and learners of the Division of Negros Oriental. The contents of this module are carefully researched, chosen, and
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13 NegOr_Q2_GenChem1-12_SLK Week1_v2
SYNOPSIS ANSWER KEY

This Self-Learning Kit deals with

the different type quantum
numbers, its meaning, and its
guiding principles. Learning
quantum numbers will make it easy
to understand the placement of
electron in an atom and at the same
time writing the electronic
configuration. In addition, based on
the notation of the electronic
configuration will explain the
paramagnetic and diamagnetic
property of an atom or an element.
So, join us, as we dive in
through the world of electronic
configuration and orbital diagram.
Pre-activity/Pre-test
I. 1. Energy level
2. Paulis exclusion principle
3. Spin quantum number
4. Neils Bohr
ABOUT THE AUTHOR 5. Valence electrons
6. Aufbau principle
MARIA CRISTINA M. SALASALAN
7. d orbitalsis a graduate of Bachelor
of Science in Nursing 8.
from
Werner Jose Rizal Memorial State
Heisenberg
9. Subshells
University in 2010. Finished
10. 1
Continuing Professional
Education (CPE) at Villaflores College in 2015. Finished her
Complete Academic Requirement
Evaluation/Post-test: for Master of Arts in
Science Teaching at Negros Oriental
I. 1. a 2. b 3. c 4. a 5. State
c University.
II. 1. Be = 1s2 2s22p1
Currently working at Department of
2. Si = 1s2 2s2 3p6 3s2 3p2
Education, as
teacher II, connected at3.Crisostomo
H = 1s1 O. Retes National
High School Senior High. 4. Be = 1s 2 2s2

5. Cl = 1s2 2s2 2p6 3s2 3p5

III.

1. = paramagnetic

2. = diamagnetic

3.
= paramagnetic

14 NegOr_Q2_GenChem1-12_SLK Week1_v2