Electron orbital: Difference between revisions

Latest revision as of 06:01, 11 August 2024

Main Article

Discussion

Definitions of orbitals

Several kinds of orbitals can be distinguished.

Atomic orbital (AO)

See the article atomic orbital

Molecular orbital (MO)

See the article molecular orbital

Spinorbitals

The AOs and MOs defined so far depend only on the spatial coordinate vector r of a single electron. In addition, an electron has a spin coordinate μ, which can have two values: spin-up or spin-down. A complete set of functions of μ consists of two functions only, traditionally these are denoted by α(μ) and β(μ). These functions are eigenfunctions of the z-component s_z of the spin angular momentum operator with eigenvalues ±½.

Spin atomic orbital

The most general spin atomic orbital of an electron is of the form

\chi _{i}^{+}(\mathbf {r} )\alpha (\mu )+\chi _{i}^{-}(\mathbf {r} )\beta (\mu ),

where r is a vector from the nucleus of the atom to the position of the electron. In general this function is not an eigenfunction of s_z. More common is the use of either

\chi _{i}(\mathbf {r} )\alpha (\mu )\quad {\hbox{or}}\quad \chi _{i}(\mathbf {r} )\beta (\mu ),

which are eigenfunctions of s_z. Since it is rare that different AOs are used for spin-up and spin-down electrons, we dropped the superscripts + and −.

Spin molecular orbital

A spin molecular orbital is usually either

\phi ^{+}(\mathbf {r} )\alpha (\mu )\quad {\hbox{or}}\quad \phi ^{-}(\mathbf {r} )\beta (\mu ).

Here the superscripts + and − might be necessary, because some quantum chemical methods distinguish the spatial wave functions of electrons with different spins. These are the so-called different orbitals for different spins (DODS) (or spin-unrestricted) methods. However, many quantum chemical methods apply the spin-restriction:

\phi ^{+}(\mathbf {r} )=\phi ^{-}(\mathbf {r} )=\phi (\mathbf {r} ).

Chemists express this spin-restriction by stating that two electrons [an electron with spin up (α), and another electron with spin down (β)] are placed in the same spatial orbital φ. This means that the total N-electron wave function contains a factor of the type φ(r₁)α(μ₁)×φ(r₂)β(μ₂).

Applications

Before the advent of electronic computers, orbitals (molecular as well as atomic) were used extensively in qualitative arguments explaining all kinds of properties of atoms and molecules. Orbitals still play this role in introductory texts and also in organic chemistry, where orbitals serve in the explanation of some reaction mechanisms, for instance in the Woodward-Hoffmann rules. In modern computational quantum chemistry the role of atomic orbitals is different; they serve as a convenient expansion basis, comparable to powers of x in a Taylor expansion of a function f(x), or sines and cosines in a Fourier series.

Atomic orbitals

Originally AOs were defined as approximate solutions of atomic Schrödinger equations, see the section solution of the atomic Schrödinger equation in atomic orbital for more on this. Very often atomic orbitals are depicted in polar plots. In fact, the angular parts (functions depending on the spherical polar angles θ and φ) are commonly plotted. The angular parts of atomic orbitals are functions known as spherical harmonics.

As stated, quantum chemists see AOs simply as convenient basis functions for quantum mechanical computations. See the section AO basis sets in the article atomic orbital for more on the use of atomic orbitals in quantum chemistry. See Slater orbital for the explicit analytic form of orbitals and see Gauss type orbitals for a discussion of the type of AOs that are most frequently used in computations.

Molecular orbitals

To date the most widely applied method for computing molecular orbitals is a Hartree-Fock method in which the MO is expanded in atomic orbitals. Since the term "Hartree-Fock" (HF) is in quantum chemistry almost synonymous with the term "self-consistent field" (SCF), such an MO is often referred to as an SCF-LCAO-MO.

See the section computation of molecular orbitals in molecular orbital for a simple example of a case where an MO is almost completely determined by symmetry. In more complicated cases it is necessary to solve the Hartree-Fock-Roothaan equations, which have the form of a generalized eigenvalue problem. The dimension of this problem is equal to the number of atomic orbitals included in the AO basis (see AO basis sets in atomic orbital for more details).

History of term orbital

Orbit is an old noun introduced by Johannes Kepler in 1609 to describe the trajectories of the earth and the planets. The adjective "orbital" had (and still has) the meaning "relating to an orbit". When Ernest Rutherford in 1911 postulated his planetary model of the atom (the nucleus as the sun, and the electrons as the planets) it was natural to call the paths of the electrons "orbits". Bohr used the word as well, although he was the first to recognize (1913) that an electron orbit is not a trajectory, but a stationary state of the hydrogen atom.

After Schrödinger (1926) had solved his wave equation for the hydrogen atom (see hydrogen-like atoms for details), it became clear that the electronic orbits did not resemble planetary orbits at all. The wave functions of the hydrogen electron are time-independent and smeared out; they are more like unmoving clouds than planetary orbits. As a matter of fact, the angular parts of the hydrogen wave functions are spherical harmonic functions and hence they have the same appearance as these functions. (See spherical harmonics for a few graphical illustrations).

In the 1920s electron spin was discovered, whereupon the adjective "orbital" started to be used in the meaning of "non-spin", that is, as a synonym of "spatial". In scientific papers of around 1930 one finds discussions about "orbital degeneracy", meaning that the spatial (non-spin) parts of several one-electron wave functions have the same energy. Also the terms orbital- and spin-angular momentum date form these days.

In 1932 Robert S. Mulliken coined the noun "orbital". Mulliken wrote:^[2]

From here on, one-electron orbital wave functions will be referred to for brevity as orbitals.

Note that here, evidently, Mulliken uses "orbital" as relating to "spatial" (non-spin) and defines an orbital as what is now called a "spatial orbital". Then Mulliken went on in the same article to distinguish atomic and molecular orbitals.

Later the somewhat unfortunate term "spinorbital" was introduced for the product functions φ(r)α(μ) and φ(r)β(μ) in which φ(r) has the tautological name "spatial orbital" and α(μ) and β(μ) are called "one-electron spin functions". The term "spinorbital" is unfortunate because it merges in one word the concepts of spin and orbital, which were distinguished carefully by early writers on quantum mechanics. For instance, one of the pioneers of theoretical chemistry, Walter Heitler, juxtaposes two-electron spin functions and two-electron orbital functions.^[3] In the phrase "two-electron orbital function", Heitler uses orbital as an adjective synonymous with spatial (non-spin). Note, parenthetically, that Heitler does not refer to a "two-electron orbital", (there is no such thing as a two-electron orbital !) and that an inexperienced reader may easily—and erroneously—interpret the term "two-electron orbital function" as "two-electron orbital".

References and notes

↑ Here "orbital" is used as a noun. In quantum mechanics, the adjective orbital is often used as a synonym of "spatial" (as in orbital angular momentum), in contrast to spin (as in spin angular momentum).
↑ R. S. Mulliken, Electronic Structures of Molecules and Valence. II General Considerations, Physical Review, vol. 41, pp. 49-71 (1932)
↑ W. Heitler, Elementary Wave Mechanics, 2nd edition (1956) Clarendon Press, Oxford, UK.

[1] Here "orbital" is used as a noun. In quantum mechanics, the adjective orbital is often used as a synonym of "spatial" (as in orbital angular momentum), in contrast to spin (as in spin angular momentum).

[2] R. S. Mulliken, Electronic Structures of Molecules and Valence. II General Considerations, Physical Review, vol. 41, pp. 49-71 (1932)

[3] W. Heitler, Elementary Wave Mechanics, 2nd edition (1956) Clarendon Press, Oxford, UK.

[1]

[2]

[3]

@@ Line 1: / Line 1: @@
 {{subpages}}
-In [[quantum chemistry]], an '''electron orbital''' (or more often just '''orbital''') is a synonym for a [[quadratically integrable]]  one-electron [[wave function]]. Here "orbital" is used as a noun. In [[quantum mechanics]], the ''adjective'' orbital is often used as a synonym of "spatial", in contradistinction to [[spin]].
+In [[quantum chemistry]], an '''electron orbital''' (or more often just '''orbital''') is a synonym for a  one-electron function, i.e., a function of a single vector '''r''', the position vector of the electron.<ref> Here "orbital" is used as a ''noun''. In [[quantum mechanics]], the ''adjective'' orbital is often used as a synonym of "spatial" (as in orbital [[angular momentum]]), in contrast to [[spin]] (as in spin angular momentum).</ref>
+The majority of quantum chemical methods expect that an orbital has a finite [[norm]], i.e., that the orbital is normalizable (quadratically integrable), and hence this requirement is often added to the definition.
+In other branches of chemistry, an orbital is often seen as a [[wave function]] of an electron, meaning that an orbital is seen as a solution of an (effective) ''one-electron [[Schrödinger equation]]''. This point of view is a narrowing of the more general quantum chemical definition, but not contradictory to it.  In the past quantum chemists, too, distinguished one-electron functions from orbitals; one-electron functions were fairly arbitrary functions of the position vector '''r''', while  orbitals were solutions of certain effective one-electron Schrödinger equations. This distinction faded out in quantum chemistry, resulting in the present definition  of "normalizable one-electron function", not necessarily eigenfunctions of a one-electron [[Hamiltonian]].
+Usually one distinguishes two kinds of orbitals:  ''atomic orbitals'' and ''molecular orbitals''.
+Atomic orbitals are expressed with respect to one Cartesian system of axes centered on a single atom. Molecular orbitals (MOs) are "spread out" over a molecule. Usually this is a consequence of an MO being a linear combination (weighted sum) of atomic orbitals centered on different atoms.
 ==Definitions of orbitals==
 Several kinds of orbitals can be distinguished.
-===Atomic orbital===
+===Atomic orbital (AO)===
-The basic kind of orbital  is the ''atomic orbital'' (AO). This is a function depending on a single 3-dimensional vector '''r'''<sub>''A''1</sub>, which is a vector pointing from point ''A'' to electron 1.  Generally there is a nucleus at ''A''.<ref>Floating AOs and bond functions,  both of which have an empty point ''A'', are sometimes used.</ref>
+''See the article [[atomic orbital]]''
-The following notation for an AO is frequently used,
-:<math>
- \chi_i(\mathbf{r}_{A1}),
-</math>
-but other notations can be found in the literature.  Sometimes the center ''A'' is added as an index: &chi;<sub>''A i''</sub>.  We say that &chi;<sub>''A i''</sub> (or, as the case may be, &chi;<sub>'' i''</sub>) is ''centered at'' ''A''. In numerical computations AOs are either taken as [[Slater orbital| Slater type orbitals]] (STOs) or [[Gaussian type orbitals]] (GTOs). [[Hydrogen-like atom|Hydrogen-like orbitals]] are rarely applied in numerical calculations, but form the basis of many qualitative arguments in chemical bonding and atomic spectroscopy.
-The orbital is quadratically integrable, which means that the following integral is finite,
-:<math>
-\le \int_{-\infty}^{\infty}  \int_{-\infty}^{\infty} \int_{-\infty}^{\infty}\chi_i(\mathbf{r}_{A1})^* \chi_i(\mathbf{r}_{A1})\; dx_{A1}\, dy_{A1}\, dz_{A1}\ < \infty.
-</math>
-Its integrand being  real and non-negative, the integral is real and non-negative. The integral is zero if and only if &chi;<sub>'' i''</sub> is the zero function.
-===Molecular orbital===
-The second kind of orbital is the ''molecular orbital'' (MO). Such a one-electron function depends on several vectors: '''r'''<sub>''A''1</sub>, '''r'''<sub>''B''1</sub>, '''r'''<sub>''C''1</sub>, ... where ''A'', ''B'', ''C'', ... are different points in space (usually nuclear positions). The oldest example of an MO (without use of the name MO yet) is in the work of Burrau (1927) on the single-electron ion H<sub>2</sub><sup>+</sup>. Burrau applied [[spheroidal coordinates]] (a bipolar coordinate system)  to describe the wavefunctions of the electron of H<sub>2</sub><sup>+</sup>.
-Lennard-Jones<ref>J. E.Lennard-Jones, ''The Electronic Structure of some Diatomic Molecules;;  Trans. Faraday Soc. vol '''25''', p. 668 (1929).</ref> introduced the following ''linear combination of atomic orbitals'' (LCAO) way of writing an MO  &phi;:
+===Molecular orbital (MO)===
-:<math>
+''See the article [[molecular orbital]]''
-\phi(\mathbf{r}_1) = \sum_{A=1}^{N_\mathrm{nuc}} \sum_{i=1}^{n_A} c_{iA} \chi_i(\mathbf{r}_{A1}),
-\qquad c_{iA} \in \mathbb{C},
-</math>
-where ''A'' runs over  ''N''<sub>nuc</sub> different points in space (usually ''A'' runs over all the nuclei of a molecule, hence the name molecular orbital), and ''i'' runs over  the ''n''<sub>''A''</sub> different AOs centered at ''A''. The complex coefficients ''c''<sub>'' iA''</sub> can be calculated by any of the existing effective-one-electron [[quantum chemical methods]]. Examples of such methods are the [[Hückel method]] and the [[Hartree-Fock method]].
 ===Spinorbitals===
-The AOs and MOs defined so far depend only on the ''spatial coordinate vector'' '''r'''<sub>''A''1</sub> of electron 1. In addition, an electron has a [[spin coordinate]] &mu;, which can have two values: spin-up or spin-down. A complete set of functions of &mu; consists of two functions only, traditionally these are denoted by &alpha;(&mu;) and &beta;(&mu;). These functions  are eigenfunctions of the ''z''-component ''s''<sub>''z''</sub> of the [[spin angular momentum operator]] with eigenvalues &plusmn;&frac12;.
+The AOs and MOs defined so far depend only on the ''spatial coordinate vector'' '''r''' of a single electron. In addition, an electron has a [[spin coordinate]] &mu;, which can have two values: spin-up or spin-down. A complete set of functions of &mu; consists of two functions only, traditionally these are denoted by &alpha;(&mu;) and &beta;(&mu;). These functions  are eigenfunctions of the ''z''-component ''s''<sub>''z''</sub> of the [[Angular momentum (quantum)#Spin angular momentum|spin angular momentum operator]] with eigenvalues &plusmn;&frac12;.
 ====Spin atomic orbital====
-The most general ''spin atomic orbital'' of electron 1 is of the form
+The most general ''spin atomic orbital'' of an electron  is of the form
 :<math>
-  \chi_i^{+}(\mathbf{r}_{A1})\alpha(\mu_1) + \chi_i^{-}(\mathbf{r}_{A1})\beta(\mu_1)
+  \chi_i^{+}(\mathbf{r})\alpha(\mu) + \chi_i^{-}(\mathbf{r})\beta(\mu),
 </math>
-which in general is ''not'' an eigenfunction of ''s''<sub>''z''</sub>. More common is the use of either
+where '''r''' is a vector from the nucleus of the atom to the position of the electron.
+In general this function is ''not'' an eigenfunction of ''s''<sub>''z''</sub>. More common is the use of either
 :<math>
-  \chi_i(\mathbf{r}_{A1})\alpha(\mu_1) \quad\hbox{or}\quad \chi_i(\mathbf{r}_{A1})\beta(\mu_1),
+  \chi_i(\mathbf{r})\alpha(\mu) \quad\hbox{or}\quad \chi_i(\mathbf{r})\beta(\mu),
 </math>
 which are eigenfunctions of ''s''<sub>''z''</sub>. Since it is rare that different AOs are used for spin-up and spin-down electrons, we dropped the superscripts + and &minus;.
@@ Line 42: / Line 33: @@
 A ''spin molecular orbital'' is usually either
 :<math>
-  \phi^{+}(\mathbf{r}_{1})\alpha(\mu_1) \quad\hbox{or}\quad \phi^{-}(\mathbf{r}_{1})\beta(\mu_1) .
+  \phi^{+}(\mathbf{r})\alpha(\mu) \quad\hbox{or}\quad \phi^{-}(\mathbf{r})\beta(\mu) .
 </math>
 Here the superscripts + and &minus; might be necessary, because some quantum chemical methods distinguish the spatial wave functions of electrons with different spins. These are the so-called different orbitals for different spins (DODS) (or spin-unrestricted) methods. However, many quantum chemical methods apply the spin-restriction:
 :<math>
-\phi^{+}(\mathbf{r}_{1}) = \phi^{-}(\mathbf{r}_{1}) = \phi(\mathbf{r}_{1}).
+\phi^{+}(\mathbf{r}) = \phi^{-}(\mathbf{r}) = \phi(\mathbf{r}).
 </math>
-Chemists express this spin-restriction by stating that two electrons [electron 1  with spin up (&alpha;), and electron 2 with spin down (&beta;)] are placed in the same spatial orbital &phi;. This means that the total ''N''-electron wave function contains a factor of the type &phi;('''r'''<sub>1</sub>)&alpha;(&mu;<sub>1</sub>)&phi;('''r'''<sub>2</sub>)&beta;(&mu;<sub>2</sub>).
+Chemists express this spin-restriction by stating that two electrons [an electron  with spin up (&alpha;), and another electron  with spin down (&beta;)] are placed in the same spatial orbital &phi;. This means that the total ''N''-electron wave function contains a factor of the type &phi;('''r'''<sub>1</sub>)&alpha;(&mu;<sub>1</sub>)&times;&phi;('''r'''<sub>2</sub>)&beta;(&mu;<sub>2</sub>).
-==History of term==
-Orbit is an old noun (1548), initially indicating the path of the Moon and later the paths of other heavenly bodies as well. The ''adjective'' "orbital" had (and still has) the meaning "relating to an orbit".
-When [[Ernest Rutherford]] in 1911 postulated his planetary model of the atom (the nucleus as the Sun, and the electrons as the planets) it was natural to call the paths of the electrons  "orbits". Bohr, although he was the first to recognize (1913) orbits as stationary states of the hydrogen atom, used the word as well. However, after Schrödinger (1926) had solved his wave equation for the hydrogen atom (see [[hydrogen-like atom|this article]] for details), it became clear that the electronic "orbits" did not resemble planetary orbits at all. The wave functions of the hydrogen electron are time-independent and smeared out. They are more like unmoving clouds than like planetary orbits. As a matter of fact, the angular parts of the hydrogen wave functions  are [[spherical harmonics]] and hence they have the same appearance as spherical harmonics. (See [[spherical harmonics]] for a few graphical illustrations).
-In the 1920s [[electron spin]] was discovered, whereupon the adjective "orbital" started to be used in the meaning of  "non-spin", that is, as a synonym of "spatial".  In scientific papers of around 1930 one finds discussions about "orbital degeneracy", meaning that the spatial (non-spin) parts of several one-electron wave functions have the same energy. Also the terms orbital- and spin-angular momentum date form these days.
-In 1932 [[Robert S. Mulliken]]<ref>R. S. Mulliken, ''Electronic Structures of Molecules and Valence. II General Considerations'', Physical Review, vol. '''41''', pp. 49-71 (1932) </ref> coined the ''noun'' "orbital". He wrote: ''From here on, one-electron orbital wave functions will be referred to for brevity as orbitals''.<ref>Note that here, evidently, Mulliken uses the adjective "orbital" in the meaning of "spatial" (non-spin) and defines an orbital as, what is now called  a "spatial orbital".</ref> Then he went on to distinguish atomic  and molecular orbitals.
-Later the somewhat unfortunate term "spinorbital" was introduced for the product functions &phi;('''r'''<sub>1</sub>)&alpha;(&mu;<sub>1</sub>) and &phi;('''r'''<sub>1</sub>)&beta;(&mu;<sub>1</sub>) in which &phi;('''r'''<sub>1</sub>) has  the tautological name  "spatial orbital" and &alpha;(&mu;<sub>1</sub>) and &beta;(&mu;<sub>1</sub>) are called "one-electron spin functions". The term "spinorbital" is unfortunate because it merges in one word the concepts of  spin and orbital, which were distinguished carefully by early writers on quantum mechanics.  For instance, one of the pioneers of theoretical chemistry, [[Walter Heitler]],  correctly contraposes ''two-electron spin functions'' and ''two-electron orbital functions''.<ref>W. Heitler, ''Elementary Wave Mechanics'', 2nd edition (1956) Clarendon Press, Oxford, UK.</ref> In the phrase two-electron orbital function, Heitler uses ''orbital'' as an adjective synonymous with ''spatial'' (''non-spin''). Note, parenthetically, that Heitler does not refer to a "two-electron orbital", (there is no such thing as a two-electron orbital!), but that an inexperienced reader may easily misread  the term "two-electron orbital function" as "two-electron orbital".
 ==Applications==
-Before the advent of electronic computers, orbitals (molecular as well as atomic) were used extensively in qualitative arguments explaining several properties of atoms and molecules. Orbitals still play this role in introductory texts and also in organic chemistry, where orbitals serve in the explanation of some reaction mechanisms, for instance in the [[Woodward-Hoffmann rules]].
+Before the advent of electronic computers, orbitals (molecular as well as atomic) were used extensively in qualitative arguments explaining all kinds of  properties of atoms and molecules. Orbitals still play this role in introductory texts and also in organic chemistry, where orbitals serve in the explanation of some reaction mechanisms, for instance in the [[Woodward-Hoffmann rules]].
 In modern computational quantum chemistry the role of atomic orbitals is  different; they serve as a convenient expansion basis, comparable to powers of ''x'' in a Taylor expansion of a function ''f(x)'', or sines and cosines in a Fourier series.
 ===Atomic orbitals===
-Bohr<ref>N. Bohr, Zeitschrift für Physik, vol. '''9''', p. 1 (1922)</ref> was the first to see
+Originally AOs were defined as approximate  solutions of atomic Schrödinger equations,
-how atomic orbitals form a basis for an understanding of the [[Periodic Table]] of elements. In the explanation of the Periodic Table, atomic orbitals are labeled as if they are [[hydrogen-like atom|hydrogen orbitals]], that is, by a principal quantum number ''n'' and a letter designating the [[azimuthal quantum number]] (angular momentum quantum number) ''l''. This labeling is not only correct for hydrogen-like orbitals, but also for AOs that arise as solutions from an ''N''-electron [[independent particle model]] with a central (spherically symmetric) potential field. A [[central field]] is necessary to have ''l'' as a good quantum number, and the concept "orbital" is tied to an independent particle model, such as the Hartree and the [[Hartree-Fock]] model.
+see the section ''solution of the atomic Schrödinger equation'' in [[atomic orbital]] for more on this.  Very often atomic orbitals are depicted in [[polar plots]]. In fact, the angular parts (functions depending on the [[spherical polar coordinates|spherical polar angles]] &theta; and &phi;) are commonly plotted. The angular parts of atomic orbitals are functions known as [[spherical harmonics]].
-In an independent particle model an ''Ansatz'' for the total ''N''-electron wave function is a (possibly [[antisymmetrizer|antisymmetrized]]) product of orbitals. From the effective-one-electron Schrödinger equation ensuing from this Ansatz, the actual forms of the orbitals and  orbital energies are obtained, see [[Hartree-Fock method]] for an example and more details. Two major differences between such an effective-one-electron  model and the analytic solution of the hydrogen atom are: (i) The radial functions are given numerically, and are no longer known analytic functions, such as the Laguerre functions for the hydrogen atom. The angular parts, however, are the same analytic functions as for the H-atom (spherical harmonics). (ii) The orbitals within a given atomic shell (orbitals with same pricipal quantum number ''n'') are no longer degenerate (have no longer the same energy). While in the hydrogen atom, for instance the orbitals in the ''n'' = 3 shell: 3''s'', 3''p'', and 3''d'', all have the same energy (proportional to 1/''n''<sup>2</sup> = 1/3<sup>2</sup>), the degeneracy for the ''N''-electron atom is lifted: only those of same azimuthal quantum number ''l'' are degenerate.
+As stated, quantum chemists see AOs simply  as convenient basis functions for quantum mechanical computations. See the section ''AO basis sets'' in the article [[atomic orbital]] for more on the use of atomic orbitals in quantum chemistry. See  [[Slater orbital]] for the explicit analytic form of orbitals and see [[Gauss type orbitals]] for a discussion of the type of AOs that are most frequently used in computations.
-The principal quantum number is now simply a counting index, it counts in increasing energy orbitals of the same ''l'', starting at ''n'' = ''l'' + 1 (to be in line with the hydrogen AOs).
-It is found<ref>J. C. Slater, ''Quantum Theory of Atomic Structure'', vol. I, McGraw-Hill, New York (1960), p. 193</ref> that the orbitals arising from the central field, independent particle model have the following order in increasing energy
-:<em> 1s, 2s, 2p, 3s, 3p, 3d, 4s, 4p, 4d, 5s, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d.</em>
-Remark: for some atoms the order of 3''d'' and 4''s'' is flipped and for some other atoms the order of 4''d'' and 5''s'' is flipped.
-One can now build up the atoms in the Periodic Table by the ''Aufbau'' (building-up) principle: fill orbitals in increasing energy. Doing this one must obey the Pauli exclusion priciple that forbids  more than two electrons per spatial orbital. Allowed is at most one electron with &alpha; spin  and one with &beta; spin. Recall further that there are 2''l''+1 orbitals of definite ''l''.
-For instance, the neon atom (atomic number ''Z'' = 10) has the electronic configuration (''n'' = 1 and ''n'' = 2 shells are completely filled):
-:<math>{\scriptstyle  1s^2\, 2s^2\, 2p^6,} </math>
-while chlorine (''Z'' = 17) has
-:<math>{\scriptstyle 1s^2\, 2s^2\, 2p^6\, 3s^2\, 3p^5\,}. </math>
-For more details the article [[Periodic Table]] may be consulted.
 ===Molecular orbitals===
-In the hands of [[Friedrich Hund]],  [[Robert S. Mulliken]], [[John Lennard-Jones]], and others, [[molecular orbital theory]] was established firmly in the 1930s as a (mostly qualitative)  theory explaining much of chemical bonding, especially the bonding in diatomic molecules.
+To date the most widely applied method for computing molecular orbitals is a [[Hartree-Fock]] method
-[[Image:MOdiagram H2.png|right|250px|thumb|MO energy level diagram for H<sub>2</sub>]]
+in which the MO is expanded in  [[atomic orbital]]s. Since the term "Hartree-Fock" (HF) is in quantum chemistry almost synonymous with the term "self-consistent field" (SCF), such an MO is often referred to as an SCF-LCAO-MO.
-To give the flavor of the theory we look at the simplest molecule: H<sub>2</sub>.
+See the section ''computation of molecular orbitals'' in [[molecular orbital]] for a simple example of a case where an MO is almost completely determined by symmetry. In more complicated cases it is necessary to solve the [[Hartree]]-[[Fock]]-[[Roothaan]] equations, which have the form of a generalized [[eigenvalue problem]]. The dimension of this problem is equal to the number of atomic orbitals included in the AO basis (see ''AO basis sets'' in [[atomic orbital]] for more details).
-The two hydrogen atoms, labeled ''A'' and ''B'', each have one electron (a red arrrow in the figure) in an 1''s'' atomic orbital (AO). These AOs  are labeled in the figure ''1s''<sub>A</sub> and ''1s''<sub>B</sub>.  When the atoms start to interact the AOs combine linearly to two molecular orbitals:
-:<math>
-\sigma_g = N_g (1s_A + 1s_B)\quad \hbox{and}\quad\sigma_u = N_u (1s_A - 1s_B).
-</math>
-Earlier in this article a general expression for an LCAO-MO was given. In the present example of the &sigma;<sub>''g''</sub> MO of H<sub>2</sub>, we have ''N''<sub>nuc</sub> = 2, ''n''<sub>''A''</sub> = ''n''<sub>''B''</sub> = 1, &chi;<sub>''A''1</sub> = 1''s''<sub>''A''</sub>, &chi;<sub>''B''1</sub> = 1''s''<sub>''B''</sub>,  ''c''<sub>''A''1</sub> = ''c''<sub>''B''1</sub> = ''N''<sub>''g''</sub>.
-With regard to the form of the MOs the following: Because both hydrogen atoms are identical, the molecule has reflection symmetry with respect to a mirror plane halfway the H&mdash;H bond and perpendicular to it. It is one of the basic assumptions in quantum mechanics that wave functions show the symmetry of the system. [Technically: solutions of the Schrödinger equation belong to a subspace of Hilbert (function) space that is  irreducible  under the symmetry group of the system].
-Clearly &sigma;<sub>''g''</sub> is symmetric (is even, stays the same) under
-:<math>
-\mathrm{{\scriptstyle Reflection:}} \quad  1s_A \longleftrightarrow 1s_B,
-</math>
-while &sigma;<sub>''u''</sub> changes sign (is antisymmetric, is odd). The Greek letter &sigma; indicates invariance under rotation around the bond axis. The subscripts ''g'' and ''u'' stand for the German words ''gerade'' (even) and ''ungerade'' (odd).
-So, because of the high symmetry of the molecule we can immediately write down two molecular orbitals, which evidently are linear combinations of atomic orbitals.
-The normalization factors ''N''<sub>''g''</sub> and ''N''<sub>''u''</sub> and the orbital energies are still to be computed. The normalization constants follow from requiring the MOs to be normalized to unity. In the computation we will use the [[bra-ket]] notation for the integral over ''x''<sub>''A''1</sub>, ''y''<sub>''A''1</sub>, and ''z''<sub>''A''1</sub> (or over ''x''<sub>''B''1</sub>, ''y''<sub>''B''1</sub>, and ''z''<sub>''B''1</sub> when that is easier).
+==History of term orbital==
-:<math>
+Orbit is an old noun introduced by [[Johannes Kepler]] in 1609 to describe the trajectories of the earth and the planets. The ''adjective'' "orbital" had (and still has) the meaning "relating to an orbit". When [[Ernest Rutherford]] in 1911 postulated his planetary model of the atom (the nucleus as the sun, and the electrons as the planets) it was natural to call the paths of the electrons  "orbits". Bohr used the word as well, although he was the first to recognize (1913) that an electron orbit is not a trajectory,  but a stationary state of the hydrogen atom.
-=\langle \sigma_g | \sigma_g \rangle = N^2_g\left( \langle 1s_A | 1s_A \rangle + \langle 1s_B | 1s_B \rangle +  \langle 1s_A | 1s_B \rangle + \langle 1s_B | 1s_A \rangle \right) =
-N^2_g(2 + 2 S).
-</math>
-We used here that the AOs are normalized to unity and that the overlap integral is real
-:<math>
-\langle 1s_A | 1s_B \rangle = \langle 1s_B | 1s_A \rangle \equiv S.
-</math>
-Applying the same procedure for ''N''<sub>''u''</sub> we find
-:<math>
-N_g = \left[ 2 + 2 S \right]^{-1/2}\quad\hbox{and}\quad N_u = \left[ 2 - 2 S \right]^{-1/2}.
-</math>
-At this point one often assumes that ''S'' &cong; 0, so that both normalization factors are equal to &frac12;&radic;2.
-In order to calculate the energy of an orbital we introduce an effective-one-electron energy operator (Hamiltonian) ''h'',
+After Schrödinger (1926) had solved his wave equation for the hydrogen atom (see [[hydrogen-like atom]]s for details), it became clear that the electronic orbits did not resemble planetary orbits at all. The wave functions of the hydrogen electron are time-independent and smeared out; they are more like unmoving clouds than  planetary orbits. As a matter of fact, the angular parts of the hydrogen wave functions  are spherical harmonic functions and hence they have the same appearance as these functions. (See [[spherical harmonics]] for a few graphical illustrations).
-:<math>
-\begin{align}
-E_{\sigma_g} &=
-\langle \sigma_g |h| \sigma_g \rangle \\
-& = \frac{1}{2}\left( \langle 1s_A |h| 1s_A \rangle + \langle 1s_B |h| 1s_B \rangle +  \langle 1s_A |h| 1s_B \rangle + \langle 1s_B |h| 1s_A \rangle \right)  \\
-&\equiv \frac{1}{2}(q_A + q_B + \beta_{AB}+ \beta_{BA} ) = q + \beta.
-\end{align}
-</math>
-Since ''A'' and ''B'' are identical atoms equipped with identical AOs, we were allowed to use
-:<math>
-q_A = q_B = q = \langle 1s |h| 1s \rangle = E_{1s} \quad\hbox{and}\quad  \beta_{AB} = \beta_{BA}  =  \beta.
-</math>
-Likewise
-:<math>
-E_{\sigma_u} = \langle \sigma_u |h| \sigma_u \rangle = q - \beta.
-</math>
-The energy term &beta; is negative (causes the attraction between the atoms).  One may tempted to assume that ''q'' = &minus;&frac12; hartree (the energy of the 1''s'' orbital in the free atom). This is not the case, however, because ''h'' contains the attraction with both nuclei, so that ''q'' is distance dependent.
-If we now look at the figure in which energies are increasing in vertical direction, we see that the two AOs are at the same energy level ''q'', and that the energy of  &sigma;<sub>''g''</sub>  is a distance |&beta;| below this level, while the energy of  &sigma;<sub>''u''</sub>  is a distance |&beta;| above this level. In the present simple-minded effective-one-electron model we can simple add the orbital energies and find that the bonding energy in H<sub>2</sub> is 2&beta; (two electrons in the ''bonding MO'' &sigma;<sub>''g''</sub>, ignoring the distance dependence of ''q''). This model predicts that the bonding energy  in the one-electron ion  H<sub>2</sub><sup>+</sup> is half that of H<sub>2</sub>, which is correct within a 20% margin.
+In the 1920s [[electron spin]] was discovered, whereupon the adjective "orbital" started to be used in the meaning of  "non-spin", that is, as a synonym of "spatial".  In scientific papers of around 1930 one finds discussions about "orbital degeneracy", meaning that the spatial (non-spin) parts of several one-electron wave functions have the same energy. Also the terms orbital- and spin-angular momentum date form these days.
-If we apply the model to He<sub>2</sub>, which has four electrons, we find that we must place two electrons in the bonding orbital and two electrons in the ''antibonding MO'' &sigma;<sub>''u''</sub>, with the total energy being 2&beta; &minus; 2&beta; = 0. So, this simple application of molecular orbital theory predicts that H<sub>2</sub> is bound and that He<sub>2</sub> is not, which is in agreement with the observed facts.
+In 1932 [[Robert S. Mulliken]] coined the ''noun'' "orbital". Mulliken wrote:<ref>R. S. Mulliken, ''Electronic Structures of Molecules and Valence. II General Considerations'', Physical Review, vol. '''41''', pp. 49-71 (1932) </ref>
+<blockquote>
+''From here on, one-electron orbital wave functions will be referred to for brevity as orbitals''.
+</blockquote>
+Note that here, evidently, Mulliken uses "orbital" as relating to "spatial" (non-spin) and defines an orbital as what is now called a "spatial orbital". Then Mulliken went on in the same article to distinguish atomic  and molecular orbitals.
-'''(to be continued)'''
+Later the somewhat unfortunate term "spinorbital" was introduced for the product functions &phi;('''r''')&alpha;(&mu;) and &phi;('''r''')&beta;(&mu;) in which &phi;('''r''') has  the tautological name  "spatial orbital" and &alpha;(&mu;) and &beta;(&mu;) are called "one-electron spin functions". The term "spinorbital" is unfortunate because it merges in one word the concepts of  spin and orbital, which were distinguished carefully by early writers on quantum mechanics. For instance, one of the pioneers of theoretical chemistry, [[Walter Heitler]],  juxtaposes ''two-electron spin functions'' and ''two-electron orbital functions''.<ref>W. Heitler, ''Elementary Wave Mechanics'', 2nd edition (1956) Clarendon Press, Oxford, UK.</ref> In the phrase "two-electron orbital function", Heitler uses ''orbital'' as an adjective synonymous with ''spatial'' (''non-spin''). Note, parenthetically, that Heitler does not refer to a "two-electron orbital", (there is no such thing as a two-electron orbital !) and that an inexperienced reader may easily&mdash;and erroneously&mdash;interpret the term "two-electron orbital function" as "two-electron orbital".
 ==References and notes==
-<references />
+<references />[[Category:Suggestion Bot Tag]]
-'''(To be continued)'''

Electron orbital: Difference between revisions

Latest revision as of 06:01, 11 August 2024

Contents

Definitions of orbitals

Atomic orbital (AO)

Molecular orbital (MO)

Spinorbitals

Spin atomic orbital

Spin molecular orbital

Applications

Atomic orbitals

Molecular orbitals

History of term orbital

References and notes

Navigation menu

Electron orbital: Difference between revisions

Latest revision as of 06:01, 11 August 2024

Definitions of orbitals

Atomic orbital (AO)

Molecular orbital (MO)

Spinorbitals

Spin atomic orbital

Spin molecular orbital

Applications

Atomic orbitals

Molecular orbitals

History of term orbital

References and notes

Navigation menu

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