A Closer Look at Shells, Subshells, and OrbitalsSubshellsProblems

A total of four quantum numbers are used to describe completely the movement and trajectories of each electron within an atom. The combination of all quantum numbers of all electrons in an atom is described by a wave function that complies with the Schrödinger equation. Each electron in an atom has a unique set of quantum numbers; according to the Pauli Exclusion Principle, no two electrons can share the same combination of four quantum numbers. Quantum numbers are important because they can be used to determine the electron configuration of an atom and the probable location of the atom"s electrons. Quantum numbers are also used to understand other characteristics of atoms, such as ionization energy and the atomic radius.

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In atoms, there are a total of four quantum numbers: the principal quantum number (n), the orbital angular momentum quantum number (l), the magnetic quantum number (ml), and the electron spin quantum number (ms). The principal quantum number, \(n\), describes the energy of an electron and the most probable distance of the electron from the nucleus. In other words, it refers to the size of the orbital and the energy level an electron is placed in. The number of subshells, or \(l\), describes the shape of the orbital. It can also be used to determine the number of angular nodes. The magnetic quantum number, ml, describes the energy levels in a subshell, and ms refers to the spin on the electron, which can either be up or down.


The Principal Quantum Number (\(n\))

The principal quantum number, \(n\), designates the principal electron shell. Because n describes the most probable distance of the electrons from the nucleus, the larger the number n is, the farther the electron is from the nucleus, the larger the size of the orbital, and the larger the atom is. n can be any positive integer starting at 1, as \(n=1\) designates the first principal shell (the innermost shell). The first principal shell is also called the ground state, or lowest energy state. This explains why \(n\) can not be 0 or any negative integer, because there exists no atoms with zero or a negative amount of energy levels/principal shells. When an electron is in an excited state or it gains energy, it may jump to the second principle shell, where \(n=2\). This is called absorption because the electron is "absorbing" photons, or energy. Known as emission, electrons can also "emit" energy as they jump to lower principle shells, where n decreases by whole numbers. As the energy of the electron increases, so does the principal quantum number, e.g., n = 3 indicates the third principal shell, n = 4 indicates the fourth principal shell, and so on.

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Example \(\PageIndex{1}\)

If n = 7, what is the principal electron shell?


Example \(\PageIndex{2}\)

If an electron jumped from energy level n = 5 to energy level n = 3, did absorption or emission of a photon occur?

Answer

Emission, because energy is lost by release of a photon.


The Orbital Angular Momentum Quantum Number (\(l\))

The orbital angular momentum quantum number \(l\) determines the shape of an orbital, and therefore the angular distribution. The number of angular nodes is equal to the value of the angular momentum quantum number \(l\). (For more information about angular nodes, see Electronic Orbitals.) Each value of \(l\) indicates a specific s, p, d, f subshell (each unique in shape.) The value of \(l\) is dependent on the principal quantum number \(n\). Unlike \(n\), the value of \(l\) can be zero. It can also be a positive integer, but it cannot be larger than one less than the principal quantum number (\(n-1\)):

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Example \(\PageIndex{3}\)

If \(n = 7\), what are the possible values of \(l\)?

Answer

Since \(l\) can be zero or a positive integer less than (\(n-1\)), it can have a value of 0, 1, 2, 3, 4, 5 or 6.


Example \(\PageIndex{4}\)

If \(l = 4\), how many angular nodes does the atom have?

Answer

The number of angular nodes is equal to the value of l, so the number of nodes is also 4.


The Magnetic Quantum Number (\(m_l\))

The magnetic quantum number \(m_l\) determines the number of orbitals and their orientation within a subshell. Consequently, its value depends on the orbital angular momentum quantum number \(l\). Given a certain \(l\), \(m_l\) is an interval ranging from \(–l\) to \(+l\), so it can be zero, a negative integer, or a positive integer.

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Example \(\PageIndex{5}\)

Example: If \(n=3\), and \(l=2\), then what are the possible values of \(m_l\)?

Answer

Since \(m_l\) must range from \(–l\) to \(+l\), then \(m_l\) can be: -2, -1, 0, 1, or 2.


The Electron Spin Quantum Number (\(m_s\))

Unlike \(n\), \(l\), and \(m_l\), the electron spin quantum number \(m_s\) does not depend on another quantum number. It designates the direction of the electron spin and may have a spin of +1/2, represented by↑, or –1/2, represented by ↓. This means that when \(m_s\) is positive the electron has an upward spin, which can be referred to as "spin up." When it is negative, the electron has a downward spin, so it is "spin down." The significance of the electron spin quantum number is its determination of an atom"s ability to generate a magnetic field or not. (Electron Spin.)

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Example \(\PageIndex{5}\)

List the possible combinations of all four quantum numbers when \(n=2\), \(l=1\), and \(m_l=0\).

Answer

The fourth quantum number is independent of the first three, allowing the first three quantum numbers of two electrons to be the same. Since the spin can be +1/2 or =1/2, there are two combinations:

\(n=2\), \(l=1\), \(m_l =0\), \(m_s=+1/2\) \(n=2\), \(l=1\), \(m_l=0\), \(m_s=-1/2\)

Example \(\PageIndex{6}\)

Can an electron with \(m_s=1/2\) have a downward spin?

Answer

No, if the value of \(m_s\) is positive, the electron is "spin up."


A Closer Look at Shells, Subshells, and Orbitals


Principal Shells

The value of the principal quantum number n is the level of the principal electronic shell (principal level). All orbitals that have the same n value are in the same principal level. For example, all orbitals on the second principal level have a principal quantum number of n=2. When the value of n is higher, the number of principal electronic shells is greater. This causes a greater distance between the farthest electron and the nucleus. As a result, the size of the atom and its atomic radius increases.

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Because the atomic radius increases, the electrons are farther from the nucleus. Thus it is easier for the atom to expel an electron because the nucleus does not have as strong a pull on it, and the ionization energy decreases.



Subshells

The number of values of the orbital angular number l can also be used to identify the number of subshells in a principal electron shell:

When n = 1, l= 0 (l takes on one value and thus there can only be one subshell) When n = 2, l= 0, 1 (l takes on two values and thus there are two possible subshells) When n = 3, l= 0, 1, 2 (l takes on three values and thus there are three possible subshells)

After looking at the examples above, we see that the value of n is equal to the number of subshells in a principal electronic shell:

Principal shell with n = 1 has one subshell Principal shell with n = 2 has two subshells Principal shell with n = 3 has three subshells

To identify what type of possible subshells n has, these subshells have been assigned letter names. The value of l determines the name of the subshell:

Name of Subshell Value of \(l\)
s subshell 0
p subshell 1
d subshell 2
f subshell 3

Therefore:

Principal shell with n = 1 has one s subshell (l = 0) Principal shell with n = 2 has one s subshell and one p subshell (l = 0, 1) Principal shell with n = 3 has one s subshell, one p subshell, and one d subshell (l = 0, 1, 2)

We can designate a principal quantum number, n, and a certain subshell by combining the value of n and the name of the subshell (which can be found using l). For example, 3p refers to the third principal quantum number (n=3) and the p subshell (l=1).

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Orbitals

The number of orbitals in a subshell is equivalent to the number of values the magnetic quantum number ml takes on. A helpful equation to determine the number of orbitals in a subshell is 2l +1. This equation will not give you the value of ml, but the number of possible values that ml can take on in a particular orbital. For example, if l=1 and ml can have values -1, 0, or +1, the value of 2l+1 will be three and there will be three different orbitals. The names of the orbitals are named after the subshells they are found in:

s orbitalsp orbitalsd orbitalsf orbitals
l 0 1 2 3
ml 0 -1, 0, +1 -2, -1, 0, +1, +2 -3, -2, -1, 0, +1, +2, +3
Number of orbitals in designated subshell 1 3 5 7

In the figure below, we see examples of two orbitals: the p orbital (blue) and the s orbital (red). The red s orbital is a 1s orbital. To picture a 2s orbital, imagine a layer similar to a cross section of a jawbreaker around the circle. The layers are depicting the atoms angular nodes. To picture a 3s orbital, imagine another layer around the circle, and so on and so on. The p orbital is similar to the shape of a dumbbell, with its orientation within a subshell depending on ml. The shape and orientation of an orbital depends on l and ml.

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To visualize and organize the first three quantum numbers, we can think of them as constituents of a house. In the following image, the roof represents the principal quantum number n, each level represents a subshell l, and each room represents the different orbitals ml in each subshell. The s orbital, because the value of ml can only be 0, can only exist in one plane. The p orbital, however, has three possible values of ml and so it has three possible orientations of the orbitals, shown by Px, Py, and Pz. The pattern continues, with the d orbital containing 5 possible orbital orientations, and f has 7:

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Principle Quantum Number 4.jpg