Azimuthal quantum number

In quantum mechanics, the azimuthal quantum number is a quantum number for an atomic orbital that determines its orbital angular momentum and describes the shape of the orbital. The azimuthal quantum number is the second of a set of quantum numbers that describe the unique quantum state of an electron (the others being the principal quantum number $n$ , the magnetic quantum number $m ℓ$ , and the spin quantum number $m s$ ). It is also known as the orbital angular momentum quantum number, orbital quantum number, subsidiary quantum number, or second quantum number, and is symbolized as $ℓ$ (pronounced ell).

Derivation

There are four quantum numbers—n, ℓ, m_ℓ, m_s— connected with the energy states of an atom's electrons. These four numbers specify the unique and complete quantum state of any single electron in the atom, and they combine to compose the electron's wavefunction, or orbital.

When solving to obtain the wave function, the Schrödinger equation resolves into three equations that lead to the first three quantum numbers, meaning that the three equations are interrelated. The azimuthal quantum number arises in solving the polar part of the wave equation—relying on the spherical coordinate system, which generally works best with models having sufficient aspects of spherical symmetry.

An electron's angular momentum, $L$ , is related to its quantum number $ℓ$ by the following equation: $\mathbf {L} ^{2}\Psi =\hbar ^{2}\ell (\ell +1)\Psi ,$ where $ħ$ is the reduced Planck's constant, $L$ is the orbital angular momentum operator and $\Psi$ is the wavefunction of the electron. The quantum number $ℓ$ is always a non-negative integer: 0, 1, 2, 3, etc. (Notably, $L$ has no real meaning except in its use as the angular momentum operator; thus, it is standard practice to use the quantum number $ℓ$ when referring to angular momentum).

Atomic orbitals have distinctive shapes, (see top graphic) in which letters, s, p, d, f, etc, (employing a convention originating in spectroscopy) denote the shape of the atomic orbital. The wavefunctions of these orbitals take the form of spherical harmonics, and so are described by Legendre polynomials. The several orbitals relating to the different (integer) values of ℓ are sometimes called sub-shells—referred to by lowercase Latin letters chosen for historical reasons—as shown in the table "Quantum subshells for the azimuthal quantum number".

Quantum subshells for the azimuthal quantum number
Azimuthal quantum number (ℓ)	Historical letter	Historical name	Maximum electrons	Shape
0	s	sharp	2	Spherical atomic orbital; (see: this picture of spherical harmonics, top row).
1	p	principal	6	Three dumbbell-shaped polar-aligned orbitals; one lobe on each pole of the x, y, and z axes (on both + and − axes).
2	d	diffuse	10	Nine dumbbells and one doughnut, or "Unique shape #1"; (see: this picture of spherical harmonics, third row center).
3	f	fundamental	14	"Unique shape #2"; (see: this picture of spherical harmonics, bottom row center).
4	g	gallant	18
5	h		22
6	i		26
The letters after the g sub-shell follow in alphabetical order—excepting letter j and those already used.

Each of the different angular momentum states can take 2(2ℓ + 1) electrons. This is because the third quantum number m_ℓ (which can be thought of loosely as the quantized projection of the angular momentum vector on the z-axis) runs from −ℓ to ℓ in integer units, and so there are 2ℓ + 1 possible states. Each distinct n, ℓ, m_ℓ orbital can be occupied by two electrons with opposing spins (given by the quantum number m_s = ±1⁄2), giving 2(2ℓ + 1) electrons overall. Orbitals with higher ℓ than given in the table are perfectly permissible, but these values cover all atoms so far discovered.

For a given value of the principal quantum number n, the possible values of ℓ range from 0 to $n - 1$ ; therefore, the $n = 1$ shell only possesses an s subshell and can only take 2 electrons, the $n = 2$ shell possesses an s and a p subshell and can take 8 electrons overall, the $n = 3$ shell possesses s, p, and d subshells and has a maximum of 18 electrons, and so on.

A simplistic one-electron model results in energy levels depending on the principal number alone. In more complex atoms these energy levels split for all $n > 1$ , placing states of higher ℓ above states of lower ℓ. For example, the energy of 2p is higher than of 2s, 3d occurs higher than 3p, which in turn is above 3s, etc. This effect eventually forms the block structure of the periodic table. No known atom possesses an electron having ℓ higher than three (f) in its ground state.

The angular momentum quantum number, ℓ, governs^[how?] the number of planar nodes going through the nucleus. A planar node can be described in an electromagnetic wave as the midpoint between crest and trough, which has zero magnitudes. In an s orbital, no nodes go through the nucleus, therefore the corresponding azimuthal quantum number ℓ takes the value of 0. In a p orbital, one node traverses the nucleus and therefore ℓ has the value of 1. $L$ has the value ${\sqrt {2}}\hbar$ .

Depending on the value of n, there is an angular momentum quantum number ℓ and the following series. The wavelengths listed are for a hydrogen atom:

$n=1,L=0$ , Lyman series (ultraviolet)
$n=2,L={\sqrt {2}}\hbar$ , Balmer series (visible)
$n=3,L={\sqrt {6}}\hbar$ , Ritz–Paschen series (near infrared)
$n=4,L=2{\sqrt {3}}\hbar$ , Brackett series (short-wavelength infrared)
$n=5,L=2{\sqrt {5}}\hbar$ , Pfund series (mid-wavelength infrared).

Addition of quantized angular momenta

Given a quantized total angular momentum $\mathbf {j}$ which is the sum of two individual quantized angular momenta ${\boldsymbol {\ell }}_{1}$ and ${\boldsymbol {\ell }}_{2}$ , $\mathbf {j} ={\boldsymbol {\ell }}_{1}+{\boldsymbol {\ell }}_{2}$ the quantum number $j$ associated with its magnitude can range from $|\ell _{1}-\ell _{2}|$ to $\ell _{1}+\ell _{2}$ in integer steps where $\ell _{1}$ and $\ell _{2}$ are quantum numbers corresponding to the magnitudes of the individual angular momenta.

Total angular momentum of an electron in the atom

Due to the spin–orbit interaction in the atom, the orbital angular momentum no longer commutes with the Hamiltonian, nor does the spin. These therefore change over time. However the total angular momentum $J$ does commute with the one-electron Hamiltonian and so is constant. $J$ is defined through $\mathbf {J} =\mathbf {L} +\mathbf {S}$ $L$ being the orbital angular momentum and $S$ the spin. The total angular momentum satisfies the same commutation relations as orbital angular momentum, namely $[J_{i},J_{j}]=i\hbar \varepsilon _{ijk}J_{k}$ from which follows $\left[J_{i},J^{2}\right]=0$ where $J i$ stand for $J x$ , $J y$ , and $J z$ .

The quantum numbers describing the system, which are constant over time, are now $j$ and $m j$ , defined through the action of $J$ on the wavefunction $\Psi$ ${\begin{aligned}\mathbf {J} ^{2}\Psi &=\hbar ^{2}j(j+1)\Psi \\[1ex]\mathbf {J} _{z}\Psi &=\hbar m_{j}\Psi \end{aligned}}$

So that $j$ is related to the norm of the total angular momentum and $m j$ to its projection along a specified axis. The j number has a particular importance for relativistic quantum chemistry, often featuring in subscript in electron configuration of superheavy elements.

As with any angular momentum in quantum mechanics, the projection of $J$ along other axes cannot be co-defined with $J z$ , because they do not commute. The eigenvectors of $j$ , $s$ , $m j$ and parity, which are also eigenvectors of the Hamiltonian, are linear combinations of the eigenvectors of $ℓ$ , $s$ , $m ℓ$ and $m s$ .

Beyond isolated atoms

The angular momentum quantum numbers strictly refer to isolated atoms. However, they have wider uses for atoms in solids, liquids or gases. The $ℓ m$ quantum number corresponds to specific spherical harmonics and are commonly used to describe features observed in spectroscopic methods such as X-ray photoelectron spectroscopy^[1] and electron energy loss spectroscopy.^[2] (The notation is slightly different, with X-ray notation where K,L,M are used for excitations out of electron states with $l=0,1,2$ .)

The angular momentum quantum numbers are also used when the electron states are described in methods such as Kohn–Sham density functional theory^[3] or with gaussian orbitals.^[4] For instance, in silicon the electronic properties used in semiconductor device are due to the p-like states with $l=1$ centered at each atom, while many properties of transition metal depend upon the d-like states with $l=2$ .

List of angular momentum quantum numbers

Intrinsic (or spin) angular momentum quantum number, or simply spin quantum number
orbital angular momentum quantum number (the subject of this article)
magnetic quantum number, related to the orbital momentum quantum number
total angular momentum quantum number

History

The azimuthal quantum number was carried over from the Bohr model of the atom, and was posited by Arnold Sommerfeld.^[5] The Bohr model was derived from spectroscopic analysis of the atom in combination with the Rutherford atomic model. The lowest quantum level was found to have an angular momentum of zero. Orbits with zero angular momentum were considered as oscillating charges in one dimension and so described as "pendulum" orbits, but were not found in nature.^[6] In three-dimensions the orbits become spherical without any nodes crossing the nucleus, similar (in the lowest-energy state) to a skipping rope that oscillates in one large circle.

References

^ Hüfner, Stefan (2003). Photoelectron Spectroscopy. Advanced Texts in Physics. Berlin, Heidelberg: Springer Berlin Heidelberg. doi:10.1007/978-3-662-09280-4. ISBN 978-3-642-07520-9.
^ Egerton, R.F. (2011). Electron Energy-Loss Spectroscopy in the Electron Microscope. Boston, MA: Springer US. doi:10.1007/978-1-4419-9583-4. ISBN 978-1-4419-9582-7.
^ Kohn, W.; Sham, L. J. (1965). "Self-Consistent Equations Including Exchange and Correlation Effects". Physical Review. 140 (4A): A1133–A1138. doi:10.1103/PhysRev.140.A1133. ISSN 0031-899X.
^ Gill, Peter M.W. (1994), "Molecular integrals Over Gaussian Basis Functions", Advances in Quantum Chemistry, vol. 25, Elsevier, pp. 141–205, doi:10.1016/s0065-3276(08)60019-2, ISBN 978-0-12-034825-1, retrieved 2024-02-20
^ Eisberg, Robert (1974). Quantum Physics of Atoms, Molecules, Solids, Nuclei and Particles. New York: John Wiley & Sons Inc. pp. 114–117. ISBN 978-0-471-23464-7.
^ R.B. Lindsay (1927). "Note on "pendulum" orbits in atomic models". Proc. Natl. Acad. Sci. 13 (6): 413–419. Bibcode:1927PNAS...13..413L. doi:10.1073/pnas.13.6.413. PMC 1085028. PMID 16587189.

External links

[1] Hüfner, Stefan (2003). Photoelectron Spectroscopy. Advanced Texts in Physics. Berlin, Heidelberg: Springer Berlin Heidelberg. doi:10.1007/978-3-662-09280-4. ISBN 978-3-642-07520-9.

[2] Egerton, R.F. (2011). Electron Energy-Loss Spectroscopy in the Electron Microscope. Boston, MA: Springer US. doi:10.1007/978-1-4419-9583-4. ISBN 978-1-4419-9582-7.

[3] Kohn, W.; Sham, L. J. (1965). "Self-Consistent Equations Including Exchange and Correlation Effects". Physical Review. 140 (4A): A1133–A1138. doi:10.1103/PhysRev.140.A1133. ISSN 0031-899X.

[4] Gill, Peter M.W. (1994), "Molecular integrals Over Gaussian Basis Functions", Advances in Quantum Chemistry, vol. 25, Elsevier, pp. 141–205, doi:10.1016/s0065-3276(08)60019-2, ISBN 978-0-12-034825-1, retrieved 2024-02-20

[5] Eisberg, Robert (1974). Quantum Physics of Atoms, Molecules, Solids, Nuclei and Particles. New York: John Wiley & Sons Inc. pp. 114–117. ISBN 978-0-471-23464-7.

[6] R.B. Lindsay (1927). "Note on "pendulum" orbits in atomic models". Proc. Natl. Acad. Sci. 13 (6): 413–419. Bibcode:1927PNAS...13..413L. doi:10.1073/pnas.13.6.413. PMC 1085028. PMID 16587189.

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v t e Electron configuration
Electron shell Atomic orbital Quantum mechanics Introduction to quantum mechanics
Quantum numbers	Principal quantum number ( $n$ ) Azimuthal quantum number ( $ℓ$ ) Magnetic quantum number ( $m$ ) Spin quantum number ( $s$ )
Ground-state configurations	Periodic table (electron configurations) Electron configurations of the elements (data page)
Electron filling	Pauli exclusion principle Hund's rule Aufbau principle
Electron pairing	Electron pair Unpaired electron
Bonding participation	Valence electron Core electron
Electron counting rules	Octet rule 18-electron rule