Physics:Spin quantum number
In physics, the spin quantum number is a quantum number (designated s) that describes the intrinsic angular momentum (or spin angular momentum, or simply spin) of an electron or other particle. It has the same value for all particles of the same type, such as s = 1/2 for all electrons. It is an integer for all bosons, such as photons, and a half-odd-integer for all fermions, such as electrons and protons. The component of the spin along a specified axis is given by the spin magnetic quantum number, conventionally written ms.[1] The value of ms is the component of spin angular momentum, in units of the reduced Planck constant ħ, parallel to a given direction (conventionally labelled the z–axis). It can take values ranging from +s to −s in integer increments. For an electron, ms can be either ++1/2 or −+1/2 .
The phrase spin quantum number was originally used to describe the fourth of a set of quantum numbers (the principal quantum number n, the azimuthal quantum number ℓ, the magnetic quantum number m, and the spin magnetic quantum number ms), which completely describe the quantum state of an electron in an atom. Some introductory chemistry textbooks describe ms as the spin quantum number,[2][3] and s is not mentioned since its value 1/2 is a fixed property of the electron, sometimes using the variable s in place of ms.[4] Some authors discourage this usage as it causes confusion.[4] At a more advanced level where quantum mechanical operators or coupled spins are introduced, s is referred to as the spin quantum number, and ms is described as the spin magnetic quantum number[5] or as the z-component of spin sz.[6]
Spin quantum numbers apply also to systems of coupled spins, such as atoms that may contain more than one electron. Capitalized symbols are used: S for the total electronic spin, and mS or MS for the z-axis component. A pair of electrons in a spin singlet state has S = 0, and a pair in the triplet state has S = 1, with mS = −1, 0, or +1. Nuclear-spin quantum numbers are conventionally written I for spin, and mI or MI for the z-axis component.
The name "spin" comes from a geometrical spinning of the electron about an axis, as proposed by Uhlenbeck and Goudsmit. However, this simplistic picture was quickly realized to be physically unrealistic, because it would require the electrons to rotate faster than the speed of light.[7] It was therefore replaced by a more abstract quantum-mechanical description.
Magnetic nature of atoms and molecules
The spin quantum number helps to explain the magnetic properties of atoms and molecules. A spinning electron behaves like a micromagnet with a definite magnetic moment. If an atomic or molecular orbital contains two electrons, then their magnetic moments oppose and cancel each other.
If all orbitals are doubly occupied by electrons, the net magnetic moment is zero and the substance behaves as diamagnetic; it is repelled by the external magnetic field. If some orbitals are half filled (singly occupied), the substance has a net magnetic moment and is paramagnetic; it is attracted by the external magnetic field.
History
Early attempts to explain the behavior of electrons in atoms focused on solving the Schrödinger wave equation for the hydrogen atom, the simplest possible case, with a single electron bound to the atomic nucleus. This was successful in explaining many features of atomic spectra.
The solutions required each possible state of the electron to be described by three "quantum numbers". These were identified as, respectively, the electron "shell" number n, the "orbital" number ℓ, and the "orbital angular momentum" number m. Angular momentum is a so-called "classical" concept measuring the momentum[citation needed] of a mass in circular motion about a point. The shell numbers start at 1 and increase indefinitely. Each shell of number n contains n2 orbitals. Each orbital is characterized by its number ℓ, where ℓ takes integer values from 0 to n − 1, and its angular momentum number m, where m takes integer values from +ℓ to −ℓ. By means of a variety of approximations and extensions, physicists were able to extend their work on hydrogen to more complex atoms containing many electrons.
Atomic spectra measure radiation absorbed or emitted by electrons "jumping" from one "state" to another, where a state is represented by values of n, ℓ, and m. The so-called "transition rule" limits what "jumps" are possible. In general, a jump or "transition" is allowed only if all three numbers change in the process. This is because a transition will be able to cause the emission or absorption of electromagnetic radiation only if it involves a change in the electromagnetic dipole of the atom.
However, it was recognized in the early years of quantum mechanics that atomic spectra measured in an external magnetic field (see Zeeman effect) cannot be predicted with just n, ℓ, and m.
In January 1925, when Ralph Kronig was still a Columbia University Ph.D. student, he first proposed electron spin after hearing Wolfgang Pauli in Tübingen. Werner Heisenberg and Pauli immediately hated the idea: They had just ruled out all imaginable actions from quantum mechanics. Now Kronig was proposing to set the electron rotating in space. Pauli especially ridiculed the idea of spin, saying that "it is indeed very clever but of course has nothing to do with reality". Faced with such criticism, Kronig decided not to publish his theory and the idea of electron spin had to wait for others to take the credit.[8] Ralph Kronig had come up with the idea of electron spin several months before George Uhlenbeck and Samuel Goudsmit, but most textbooks credit these two Dutch physicists with the discovery.
Pauli subsequently proposed (also in 1925) a new quantum degree of freedom (or quantum number) with two possible values, in order to resolve inconsistencies between observed molecular spectra and the developing theory of quantum mechanics.
Shortly thereafter Uhlenbeck and Goudsmit identified Pauli's new degree of freedom as electron spin.
Electron spin
A spin- 1 /2 particle is characterized by an angular momentum quantum number for spin s = 1 /2. In solutions of the Schrödinger-Pauli equation, angular momentum is quantized according to this number, so that magnitude of the spin angular momentum is
[math]\displaystyle{ \| \bold{S} \| = \hbar\sqrt{s(s+1)} = \tfrac{\sqrt{3}}{2}\ \hbar ~. }[/math]
The hydrogen spectrum fine structure is observed as a doublet corresponding to two possibilities for the z-component of the angular momentum, where for any given direction z: [math]\displaystyle{ s_z = \pm \tfrac{1}{2}\hbar ~. }[/math]
whose solution has only two possible z-components for the electron. In the electron, the two different spin orientations are sometimes called "spin-up" or "spin-down".
The spin property of an electron would give rise to magnetic moment, which was a requisite for the fourth quantum number.
The magnetic moment vector of an electron spin is given by:
[math]\displaystyle{ \ \boldsymbol{\mu}_\text{s} = -\frac{e}{\ 2m\ }\ g_\text{s}\ \bold{S}\ }[/math]
where [math]\displaystyle{ -e }[/math] is the electron charge, [math]\displaystyle{ m }[/math] is the electron mass, and [math]\displaystyle{ g_\text{s} }[/math] is the electron spin g-factor, which is approximately 2.0023. Its z-axis projection is given by the spin magnetic quantum number [math]\displaystyle{ m_\text{s} }[/math] according to:
[math]\displaystyle{ \mu_z = -m_\text{s}\ g_\text{s}\ \mu_\mathsf{B} = \pm \tfrac{1}{2}\ g_\text{s}\ \mu_\mathsf{B}\ }[/math]
where [math]\displaystyle{ \ \mu_\mathsf{B}\ }[/math] is the Bohr magneton.
When atoms have even numbers of electrons the spin of each electron in each orbital has opposing orientation to that of its immediate neighbor(s). However, many atoms have an odd number of electrons or an arrangement of electrons in which there is an unequal number of "spin-up" and "spin-down" orientations. These atoms or electrons are said to have unpaired spins that are detected in electron spin resonance.
Detection of spin
When lines of the hydrogen spectrum are examined at very high resolution, they are found to be closely spaced doublets. This splitting is called fine structure, and was one of the first experimental evidences for electron spin. The direct observation of the electron's intrinsic angular momentum was achieved in the Stern–Gerlach experiment.
Stern–Gerlach experiment
The theory of spatial quantization of the spin moment of the momentum of electrons of atoms situated in the magnetic field needed to be proved experimentally. In 1922 (two years before the theoretical description of the spin was created) Otto Stern and Walter Gerlach observed it in the experiment they conducted.
Silver atoms were evaporated using an electric furnace in a vacuum. Using thin slits, the atoms were guided into a flat beam and the beam sent through an in-homogeneous magnetic field before colliding with a metallic plate. The laws of classical physics predict that the collection of condensed silver atoms on the plate should form a thin solid line in the same shape as the original beam. However, the in-homogeneous magnetic field caused the beam to split in two separate directions, creating two lines on the metallic plate.
The phenomenon can be explained with the spatial quantization of the spin moment of momentum. In atoms the electrons are paired such that one spins upward and one downward, neutralizing the effect of their spin on the action of the atom as a whole. But in the valence shell of silver atoms, there is a single electron whose spin remains unbalanced.
The unbalanced spin creates spin magnetic moment, making the electron act like a very small magnet. As the atoms pass through the in-homogeneous magnetic field, the force moment in the magnetic field influences the electron's dipole until its position matches the direction of the stronger field. The atom would then be pulled toward or away from the stronger magnetic field a specific amount, depending on the value of the valence electron's spin. When the spin of the electron is ++ 1 /2 the atom moves away from the stronger field, and when the spin is −+ 1 /2 the atom moves toward it. Thus the beam of silver atoms is split while traveling through the in-homogeneous magnetic field, according to the spin of each atom's valence electron.
In 1927 Phipps and Taylor conducted a similar experiment, using atoms of hydrogen with similar results. Later scientists conducted experiments using other atoms that have only one electron in their valence shell: (copper, gold, sodium, potassium). Every time there were two lines formed on the metallic plate.
The atomic nucleus also may have spin, but protons and neutrons are much heavier than electrons (about 1836 times), and the magnetic dipole moment is inversely proportional to the mass. So the nuclear magnetic dipole momentum is much smaller than that of the whole atom. This small magnetic dipole was later measured by Stern, Frisch and Easterman.
Electron paramagnetic resonance
For atoms or molecules with an unpaired electron, transitions in a magnetic field can also be observed in which only the spin quantum number changes, without change in the electron orbital or the other quantum numbers. This is the method of electron paramagnetic resonance (EPR) or electron spin resonance (ESR), used to study free radicals. Since only the magnetic interaction of the spin changes, the energy change is much smaller than for transitions between orbitals, and the spectra are observed in the microwave region.
Derivation
For a solution of either the nonrelativistic Pauli equation or the relativistic Dirac equation, the quantized angular momentum (see angular momentum quantum number) can be written as: [math]\displaystyle{ \Vert \mathbf{s} \Vert = \sqrt{s \, (s+1)\,} \, \hbar }[/math] where
- [math]\displaystyle{ \mathbf{s} }[/math] is the quantized spin vector or spinor
- [math]\displaystyle{ \Vert \mathbf{s}\Vert }[/math] is the norm of the spin vector
- s is the spin quantum number associated with the spin angular momentum
- [math]\displaystyle{ \hbar }[/math] is the reduced Planck constant.
Given an arbitrary direction z (usually determined by an external magnetic field) the spin z-projection is given by
- [math]\displaystyle{ s_z = m_s \, \hbar }[/math]
where ms is the secondary spin quantum number, ranging from −s to +s in steps of one. This generates 2 s + 1 different values of ms.
The allowed values for s are non-negative integers or half-integers. Fermions have half-integer values, including the electron, proton and neutron which all have s = ++ 1 /2 . Bosons such as the photon and all mesons) have integer spin values.
Algebra
The algebraic theory of spin is a carbon copy of the angular momentum in quantum mechanics theory.[9] First of all, spin satisfies the fundamental commutation relation: [math]\displaystyle{ \ [S_i, S_j ] = i\ \hbar\ \epsilon_{ijk}\ S_k\ , }[/math] [math]\displaystyle{ \ \left[S_i, S^2 \right] = 0\ }[/math] where [math]\displaystyle{ \ \epsilon_{ijk}\ }[/math] is the (antisymmetric) Levi-Civita symbol. This means that it is impossible to know two coordinates of the spin at the same time because of the restriction of the uncertainty principle.
Next, the eigenvectors of [math]\displaystyle{ \ S^2\ }[/math] and [math]\displaystyle{ \ S_z\ }[/math] satisfy: [math]\displaystyle{ \ S^2\ | s, m_s \rangle= {\hbar}^2\ s(s+1)\ | s, m_s \rangle\ }[/math] [math]\displaystyle{ \ S_z\ | s, m_s \rangle = \hbar\ m_s\ | s, m_s \rangle\ }[/math] [math]\displaystyle{ \ S_\pm\ | s, m_s \rangle = \hbar\ \sqrt{s(s+1) - m_s(m_s \pm 1)\ }\; | s, m_s \pm 1 \rangle\ }[/math] where [math]\displaystyle{ \ S_\pm = S_x \pm i S_y\ }[/math] are the ladder (or "raising" and "lowering") operators.
Energy levels from the Dirac equation
In 1928, Paul Dirac developed a relativistic wave equation, now termed the Dirac equation, which predicted the spin magnetic moment correctly, and at the same time treated the electron as a point-like particle. Solving the Dirac equation for the energy levels of an electron in the hydrogen atom, all four quantum numbers including s occurred naturally and agreed well with experiment.
Total spin of an atom or molecule
For some atoms the spins of several unpaired electrons (s1, s2, ...) are coupled to form a total spin quantum number S.[10][11] This occurs especially in light atoms (or in molecules formed only of light atoms) when spin–orbit coupling is weak compared to the coupling between spins or the coupling between orbital angular momenta, a situation known as L S coupling because L and S are constants of motion. Here L is the total orbital angular momentum quantum number.[11]
For atoms with a well-defined S, the multiplicity of a state is defined as 2S + 1. This is equal to the number of different possible values of the total (orbital plus spin) angular momentum J for a given (L, S) combination, provided that S ≤ L (the typical case). For example, if S = 1, there are three states which form a triplet. The eigenvalues of Sz for these three states are +1ħ, 0, and −1ħ.[10] The term symbol of an atomic state indicates its values of L, S, and J.
As examples, the ground states of both the oxygen atom and the dioxygen molecule have two unpaired electrons and are therefore triplet states. The atomic state is described by the term symbol 3P, and the molecular state by the term symbol 3Σ−g.
Nuclear spin
Atomic nuclei also have spins. The nuclear spin I is a fixed property of each nucleus and may be either an integer or a half-integer. The component mI of nuclear spin parallel to the z–axis can have (2I + 1) values I, I–1, ..., –I. For example, a 14N nucleus has I = 1, so that there are 3 possible orientations relative to the z–axis, corresponding to states mI = +1, 0 and −1.[12]
The spins I of different nuclei are interpreted using the nuclear shell model. Even-even nuclei with even numbers of both protons and neutrons, such as 12C and 16O, have spin zero. Odd mass number nuclei have half-integer spins, such as 3/ 2 for 7Li, 1 /2 for 13C and 5/ 2 for 17O, usually corresponding to the angular momentum of the last nucleon added. Odd-odd nuclei with odd numbers of both protons and neutrons have integer spins, such as 3 for 10B, and 1 for 14N.[13] Values of nuclear spin for a given isotope are found in the lists of isotopes for each element. (See isotopes of oxygen, isotopes of aluminium, etc. etc.)
See also
- Total angular momentum quantum number
- Rotational spectroscopy
- Basic quantum mechanics
References
- ↑ Pauling, Linus (1960). The nature of the chemical bond and the structure of molecules and crystals: an introduction to modern structural chemistry. Ithaca, N.Y.. ISBN 0-8014-0333-2. OCLC 545520.
- ↑ Petrucci, Ralph H.; Harwood, William S.; Herring, F. Geoffrey (2002). General Chemistry (8th ed.). Prentice Hall. p. 333. ISBN 0-13-014329-4.
- ↑ Whitten, Kenneth W.; Galley, Kenneth D.; Davis, Raymond E. (1992). General Chemistry (4th ed.). Saunders College Publishing. p. 196. ISBN 0-03-072373-6.
- ↑ 4.0 4.1 Perrino, Charles T.; Peterson, Donald L. (1989). "Another quantum number?". J. Chem. Educ. 66 (8): 623. doi:10.1021/ed066p623. ISSN 0021-9584. Bibcode: 1989JChEd..66..623P.
- ↑ Atkins, Peter; de Paula, Julio (2006). Atkins' Physical Chemistry (8th ed.). W.H. Freeman. p. 308. ISBN 0-7167-8759-8.
- ↑ Banwell, Colin N.; McCash, Elaine M. (1994). Fundamentals of Molecular Spectroscopy. McGraw-Hill. p. 135. ISBN 0-07-707976-0.
- ↑ Halpern, Paul (2017-11-21). "Spin: The quantum property that should have been impossible". Forbes. https://www.forbes.com/sites/startswithabang/2017/11/21/spin-the-quantum-property-that-nature-shouldnt-possess/.
- ↑ Bertolotti, Mario (2004). The History of the Laser. CRC Press. pp. 150–153. ISBN 978-1-4200-3340-3. https://books.google.com/books?id=JObDnEtzMJUC. Retrieved 22 March 2017.
- ↑ David J.Griffiths, Introduction to Quantum Mechanics, Oregon, Reed College, 2018, 166 p. ISBN:9781107189638.
- ↑ 10.0 10.1 Merzbacher, E. (1998). Quantum Mechanics (3rd ed.). John Wiley. pp. 430–431. ISBN 0-471-88702-1.
- ↑ 11.0 11.1 Atkins, P.; de Paula, J. (2006). Physical Chemistry (8th ed.). W.H. Freeman. p. 352. ISBN 0-7167-8759-8.
- ↑ Atkins, Peter; de Paula, Julio (2006). Atkins' Physical Chemistry (8th ed.). W.H. Freeman. p. 515. ISBN 0-7167-8759-8.
- ↑ Cottingham, W.N.; Greenwood, D.A. (1986). An introduction to nuclear physics. Cambridge University Press. pp. 36, 57. ISBN 0-521-31960-9.
External links
- Weiss, Michael (2001). "Full treatment of spin – including origins, evolution of spin theory, and details of the spin equations". UC Riverside. https://math.ucr.edu/home/baez/spin/spin.html.
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