This document discusses quantum numbers and their use in describing an electron's state in an atom. It introduces the four quantum numbers - principal quantum number (n), orbital angular momentum quantum number (l), magnetic quantum number (ml), and electron spin quantum number (ms). Each quantum number provides a piece of information about the electron's properties, such as energy level, orbital shape, and spin. The values of the quantum numbers are constrained by certain rules. For example, the values of ml can range from -l to +l. The document also provides examples of writing out the full set of quantum numbers for some electron configurations.

1. QUARTER
2
GENERAL CHEMISTRY 1

2. LESSON
1:
QUANTUM
NUMBERS
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3. INTRODUCTI
ON
If you ride an airplane, you need a ticket
to get in. Your ticket may likely specify a
gate number, a section number, a row,
and a seat number. No other ticket can
have the same four parts to it. It may
have the same gate, section, and seat
number, but it would have to be in a
different row. Each seat is unique and
allows only one occupant to fill it.
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4. REVIEW
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•The Bohr model is a one-
dimensional model that
uses one quantum
number to describe the
distribution of electrons in
the atom. The only
important information is
the size of the orbit, which
is described by the n
quantum number.

5. REVIEW
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•Schrodinger's model
allows the electron to
occupy three-dimensional
space. It therefore
requires three
coordinates, or three
quantum numbers to
describe the orbitals in
which electrons can be
found.

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7. QUANTUM NUMBERS
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•Specify the properties of the atomic
orbitals and the electrons in those
orbitals. An electron in an atom or
ion has four quantum numbers to
describe its state.

8. FOUR QUANTUM
NUMBERS
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•principal quantum number (n),
•the orbital angular momentum
quantum number (𝑙),
•the magnetic quantum number (𝑚𝑙),
and
•the electron spin quantum number
(𝑚𝑠)

9. FOUR QUANTUM
NUMBERS
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principal quantum number (n)
•Designates the main energy level (floor) or shell.
•Values: 1, 2, 3, …….∞. n = 7 is the highest
number of shell used for the atoms known.
•Describes the CLOUD SIZE. The larger the
value of n, the larger the cloud size.

10. FOUR QUANTUM
NUMBERS
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principal quantum number (n)
•The larger n is, the greater the average
distance of an electron in the orbital from
the nucleus, and therefore the larger (and
less stable) the orbital. The maximum
number of electrons possible in a given
shell is 2𝑛2
.

11. PRINCIPAL QUANTUM NUMBER
(N)
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12. FOUR QUANTUM
NUMBERS
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the orbital angular momentum quantum
number (𝑙)
•Designate the sub-level (apartment) where
the electron can be found.
•Give the SHAPE OF THE ORBITALS.
•Values of l: from 0 to (n-1) for each value of
n.

13. FOUR QUANTUM NUMBERS
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the orbital angular momentum
quantum number (𝑙)
•The value of l for a particular orbital is generally
designated by the letters s, p, d, and f
corresponding to l values of 0, 1, 2, and 3. for
the known atoms, only s, p, d, and f exist.

14. ANGULAR MOMENTUM QUANTUM
NUMBER
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15. FOUR QUANTUM NUMBERS
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the magnetic quantum number (𝑚𝑙)
•Designates the orbital (room) where the electron
can be found.
•Gives the DIRECTION IN SPACE that the orbital
takes.
•𝑚𝑙 specifies to which orbital within a subshell the
electron is assigned. Orbitals in a given subshell
differ only in their orientation in space, not in
their shape.

16. FOUR QUANTUM NUMBERS
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the magnetic quantum number
(𝑚𝑙)
•Values of 𝑚𝑙 : from –l, ….0, ….+l.
•The middle orbital of a subshell has a value
of 0. Orbitals to the left of the middle orbital
have negative numbers; to the right, they
have + numbers.

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18. QUESTION: IF N=3,
AND L=2, THEN WHAT
ARE THE POSSIBLE
VALUES OF 𝑚𝑙 ?
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19. FOUR QUANTUM
NUMBERS
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the electron spin quantum number
(𝑚𝑠)
•Designates the SPIN OF THE ELECTRON and
describes the behavior of the electron, not the
location.
•Values: +½, -½.
•Arrow up ↑ is +½, (referred to as “spin-up”);
arrow down ↓ is –½ (referred to as “spin down”).

20. FOUR QUANTUM
NUMBERS
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the electron spin quantum number
(𝑚𝑠)
Example 1: Determine the n, l, ml, and ms of 4𝑠1
sublevel.
Answer: n=4; l=0; ml=0; ms=+1/2

21. FOUR QUANTUM
NUMBERS
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the electron spin quantum number
(𝑚𝑠)
Example 2: Determine the n, l, ml, and ms of
3𝑝3
sublevel.
Answer: n=3; l=1; ml=+1; ms=+1/2

22. SAMPLE PROBLEM:
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23. LESSON
2:
MAGNETIC
PROPERTIES OF THE
ATOM BASED ON ITS
ELECTRONIC
CONFIGURATION
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24. INTRODUCTI
ON
Whether you see it or not, there are
thousands of chemical reactions
happening inside and outside of us
every second. Chemistry has done so
much to enable our hands to create
new things in our homes, shops, and
laboratories. Creativity and reactions
work, and it has changed the
trajectory of humanity’s way of life.
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25. REVIEW
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•Element is a substance that consists of only
one type of atom. Each atom from an element
retains all the characteristics of that element.
•Atoms are made up of subatomic particles called
protons, neutrons, and electrons and inside the
subatomic particles are elementary particles.
•The nucleus of the atom contains protons and
neutrons. Nearly all of the mass of the atom is
concentrated in the nucleus.

26. REVIEW
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• Atomic number is equal to the number of protons in the
nucleus of an atom of an element. The numbers of protons in
an atom never change. The protons are positively charged
subatomic particles.
• Neutron is an uncharged subatomic particle. When there is an
added number of neutrons in an atom of the same element,
we call them isotopes.
• The electrons are negatively charged subatomic particles
found at large distances from the nucleus. Electrons in the
outer shell of an atom are the most important in determining
chemical properties
• In the neutral state, the number of electrons is equal to the
number of protons.

27. REVIEW
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• Ion is a charged atom or substance formed by the
addition or removal of one or more electron(s) from a
neutral atom.
• A group contains elements with similar properties in
the periodic table of elements. Elements in the group
have the same number of outer electrons.
• An electronic configuration is the arrangement of
electrons in energy levels around the nucleus.

28. REVIEW
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•Magnetism is the behavior of substance
when exposed to a magnetic field. If an
element has at least one half-filled orbital or
unpaired electron, they are called
paramagnetic. If an element has no half-
filled orbital, or all electrons are paired, it is
called diamagnetic.
Reviewing these key terms is important
because foundation matters.

29. RULES AND PRINCIPLES FOR
ELECTRONIC CONFIGURATION
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1.Aufbau Principle
2. Pauli Exclusion Principle
3. The Hund’s Rule

30. 1. AUFBAU PRINCIPLE
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States that in the ground state of an atom,
electrons fill atomic orbitals of the lowest
available energy levels before occupying
higher levels.
1s > 2s > 2p >3s > 3p > 4s >3 d > 4p > 5s
> 4d > 5p > 6s > 4f . . .
The s orbital is a sphere, and it can hold 2
electrons, and it exists in all energy levels.

32. 2. PAULI EXCLUSION PRINCIPLE
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•Each atomic orbital can hold at most two
electrons,
•The two electrons must have opposite spins up-
spin (+1/2), and down-spin (-1/2)

33. 2. PAULI EXCLUSION PRINCIPLE
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34. 3. THE HUND’S RULE
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•Hund's Rule states that electrons
must occupy every orbital with a
single electron, and only then we
can put the second electron into the
orbital.

35. 3. THE HUND’S RULE
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•This may leave the atom with
many unpaired electrons.
Because unpaired electrons can
spin in either direction, they
display magnetic moments in
any direction.

36. 3. THE HUND’S RULE
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• Hund’s rule: The most stable arrangement of electrons in the subshells is the one
with the most number of parallel spins. It is a guide in determining the most
stable distribution.

38. ION IN ELECTRON
CONFIGURATION
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39. TO SUMMARIZE
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40. TO SUMMARIZE
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41. SAMPLE PROBLEM
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42. SAMPLE PROBLEM
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43. SAMPLE PROBLEM
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44. SAMPLE PROBLEM
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45. SAMPLE PROBLEM
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