Electron configuration

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Electron atomic and molecular orbitals Electron orbitals.svg
Electron atomic and molecular orbitals
A Bohr diagram of lithium Electron shell 003 Lithium - no label.svg
A Bohr diagram of lithium

In atomic physics and quantum chemistry, the electron configuration is the distribution of electrons of an atom or molecule (or other physical structure) in atomic or molecular orbitals. [1] For example, the electron configuration of the neon atom is 1s2 2s2 2p6, meaning that the 1s, 2s, and 2p subshells are occupied by two, two, and six electrons, respectively.


Electronic configurations describe each electron as moving independently in an orbital, in an average field created by the nuclei and all the other electrons. Mathematically, configurations are described by Slater determinants or configuration state functions.

According to the laws of quantum mechanics, a level of energy is associated with each electron configuration. In certain conditions, electrons are able to move from one configuration to another by the emission or absorption of a quantum of energy, in the form of a photon.

Knowledge of the electron configuration of different atoms is useful in understanding the structure of the periodic table of elements, for describing the chemical bonds that hold atoms together, and in understanding the chemical formulas of compounds and the geometries of molecules. In bulk materials, this same idea helps explain the peculiar properties of lasers and semiconductors.

Shells and subshells

s (l = 0)p (l = 1)
m = 0m = 0m = ±1
n = 1 Atomic-orbital-cloud n1 l0 m0.png
n = 2 Atomic-orbital-cloud n2 l0 m0.png Atomic-orbital-cloud n2 l1 m0.png Atomic-orbital-cloud n2 px.png Atomic-orbital-cloud n2 py.png

Electron configuration was first conceived under the Bohr model of the atom, and it is still common to speak of shells and subshells despite the advances in understanding of the quantum-mechanical nature of electrons.

An electron shell is the set of allowed states that share the same principal quantum number, n, that electrons may occupy. In each term of an electron configuration, n is the positive integer that precedes each orbital letter (helium's electron configuration is 1s2, therefore n = 1, and the orbital contains two electrons). An atom's nth electron shell can accommodate 2n2 electrons. For example, the first shell can accommodate two electrons, the second shell eight electrons, the third shell eighteen, and so on. The factor of two arises because the number of allowed states doubles with each successive shell due to electron spin—each atomic orbital admits up to two otherwise identical electrons with opposite spin, one with a spin +12 (usually denoted by an up-arrow) and one with a spin of −12 (with a down-arrow).

A subshell is the set of states defined by a common azimuthal quantum number, l, within a shell. The value of l is in the range from 0 to n  1. The values l = 0, 1, 2, 3 correspond to the s, p, d, and f labels, respectively. For example, the 3d subshell has n = 3 and l = 2. The maximum number of electrons that can be placed in a subshell is given by 2(2l + 1). This gives two electrons in an s subshell, six electrons in a p subshell, ten electrons in a d subshell and fourteen electrons in an f subshell.

The numbers of electrons that can occupy each shell and each subshell arise from the equations of quantum mechanics, [lower-alpha 1] in particular the Pauli exclusion principle, which states that no two electrons in the same atom can have the same values of the four quantum numbers. [2]


Physicists and chemists use a standard notation to indicate the electron configurations of atoms and molecules. For atoms, the notation consists of a sequence of atomic subshell labels (e.g. for phosphorus the sequence 1s, 2s, 2p, 3s, 3p) with the number of electrons assigned to each subshell placed as a superscript. For example, hydrogen has one electron in the s-orbital of the first shell, so its configuration is written 1s1. Lithium has two electrons in the 1s-subshell and one in the (higher-energy) 2s-subshell, so its configuration is written 1s2 2s1 (pronounced "one-s-two, two-s-one"). Phosphorus (atomic number 15) is as follows: 1s2 2s2 2p6 3s2 3p3.

For atoms with many electrons, this notation can become lengthy and so an abbreviated notation is used. The electron configuration can be visualized as the core electrons, equivalent to the noble gas of the preceding period, and the valence electrons: each element in a period differs only by the last few subshells. Phosphorus, for instance, is in the third period. It differs from the second-period neon, whose configuration is 1s2 2s2 2p6, only by the presence of a third shell. The portion of its configuration that is equivalent to neon is abbreviated as [Ne], allowing the configuration of phosphorus to be written as [Ne] 3s2 3p3 rather than writing out the details of the configuration of neon explicitly. This convention is useful as it is the electrons in the outermost shell that most determine the chemistry of the element.

For a given configuration, the order of writing the orbitals is not completely fixed since only the orbital occupancies have physical significance. For example, the electron configuration of the titanium ground state can be written as either [Ar] 4s2 3d2 or [Ar] 3d2 4s2. The first notation follows the order based on the Madelung rule for the configurations of neutral atoms; 4s is filled before 3d in the sequence Ar, K, Ca, Sc, Ti. The second notation groups all orbitals with the same value of n together, corresponding to the "spectroscopic" order of orbital energies that is the reverse of the order in which electrons are removed from a given atom to form positive ions; 3d is filled before 4s in the sequence Ti4+, Ti3+, Ti2+, Ti+, Ti.

The superscript 1 for a singly occupied subshell is not compulsory; for example aluminium may be written as either [Ne] 3s2 3p1 or [Ne] 3s2 3p. In atoms where a subshell is unoccupied despite higher subshells being occupied (as is the case in some ions, as well as certain neutral atoms shown to deviate from the Madelung rule), the empty subshell is either denoted with a superscript 0 or left out altogether. For example, neutral palladium may be written as either [Kr] 4d10 5s0 or simply [Kr] 4d10, and the lanthanum(III) ion may be written as either [Xe] 4f0 or simply [Xe]. [3]

It is quite common to see the letters of the orbital labels (s, p, d, f) written in an italic or slanting typeface, although the International Union of Pure and Applied Chemistry (IUPAC) recommends a normal typeface (as used here). The choice of letters originates from a now-obsolete system of categorizing spectral lines as "sharp", "principal", "diffuse" and "fundamental" (or "fine"), based on their observed fine structure: their modern usage indicates orbitals with an azimuthal quantum number, l, of 0, 1, 2 or 3 respectively. After f, the sequence continues alphabetically g, h, i... (l = 4, 5, 6...), skipping j, although orbitals of these types are rarely required. [4] [5]

The electron configurations of molecules are written in a similar way, except that molecular orbital labels are used instead of atomic orbital labels (see below).

Energy of ground state and excited states

The energy associated to an electron is that of its orbital. The energy of a configuration is often approximated as the sum of the energy of each electron, neglecting the electron-electron interactions. The configuration that corresponds to the lowest electronic energy is called the ground state. Any other configuration is an excited state.

As an example, the ground state configuration of the sodium atom is 1s2 2s2 2p6 3s1, as deduced from the Aufbau principle (see below). The first excited state is obtained by promoting a 3s electron to the 3p subshell, to obtain the 1s2 2s2 2p6 3p1 configuration, abbreviated as the 3p level. Atoms can move from one configuration to another by absorbing or emitting energy. In a sodium-vapor lamp for example, sodium atoms are excited to the 3p level by an electrical discharge, and return to the ground state by emitting yellow light of wavelength 589 nm.

Usually, the excitation of valence electrons (such as 3s for sodium) involves energies corresponding to photons of visible or ultraviolet light. The excitation of core electrons is possible, but requires much higher energies, generally corresponding to X-ray photons. This would be the case for example to excite a 2p electron of sodium to the 3s level and form the excited 1s2 2s2 2p5 3s2 configuration.

The remainder of this article deals only with the ground-state configuration, often referred to as "the" configuration of an atom or molecule.


Irving Langmuir was the first to propose in his 1919 article "The Arrangement of Electrons in Atoms and Molecules" in which, building on Gilbert N. Lewis's cubical atom theory and Walther Kossel's chemical bonding theory, he outlined his "concentric theory of atomic structure". [6] Langmuir had developed his work on electron atomic structure from other chemists as is shown in the development of the History of the periodic table and the Octet rule.

Niels Bohr (1923) incorporated Langmuir's model that the periodicity in the properties of the elements might be explained by the electronic structure of the atom. [7] His proposals were based on the then current Bohr model of the atom, in which the electron shells were orbits at a fixed distance from the nucleus. Bohr's original configurations would seem strange to a present-day chemist: sulfur was given as instead of 1s2 2s2 2p6 3s2 3p4 (2.8.6). Bohr used 4 and 6 following Alfred Werner's 1893 paper. In fact, the chemists believed in atoms long before the physicists. Langmuir began his paper referenced above by saying,

«…The problem of the structure of atoms has been attacked mainly by physicists who have given little consideration to the chemical properties which must ultimately be explained by a theory of atomic structure. The vast store of knowledge of chemical properties and relationships, such as is summarized by the Periodic Table, should serve as a better foundation for a theory of atomic structure than the relatively meager experimental data along purely physical lines... These electrons arrange themselves in a series of concentric shells, the first shell containing two electrons, while all other shells tend to hold eight.…»

The valence electrons in the atom were described by Richard Abegg in 1904. [8]

In 1924, E. C. Stoner incorporated Sommerfeld's third quantum number into the description of electron shells, and correctly predicted the shell structure of sulfur to be 2.8.6. [9] However neither Bohr's system nor Stoner's could correctly describe the changes in atomic spectra in a magnetic field (the Zeeman effect).

Bohr was well aware of this shortcoming (and others), and had written to his friend Wolfgang Pauli in 1923 to ask for his help in saving quantum theory (the system now known as "old quantum theory"). Pauli hypothesized successfully that the Zeeman effect can be explained as depending only on the response of the outermost (i.e., valence) electrons of the atom. Pauli was able to reproduce Stoner's shell structure, but with the correct structure of subshells, by his inclusion of a fourth quantum number and his exclusion principle (1925): [10]

It should be forbidden for more than one electron with the same value of the main quantum number n to have the same value for the other three quantum numbers k [l], j [ml] and m [ms].

The Schrödinger equation, published in 1926, gave three of the four quantum numbers as a direct consequence of its solution for the hydrogen atom: [lower-alpha 1] this solution yields the atomic orbitals that are shown today in textbooks of chemistry (and above). The examination of atomic spectra allowed the electron configurations of atoms to be determined experimentally, and led to an empirical rule (known as Madelung's rule (1936), [11] see below) for the order in which atomic orbitals are filled with electrons.

Atoms: Aufbau principle and Madelung rule

The aufbau principle (from the German Aufbau, "building up, construction") was an important part of Bohr's original concept of electron configuration. It may be stated as: [12]

a maximum of two electrons are put into orbitals in the order of increasing orbital energy: the lowest-energy subshells are filled before electrons are placed in higher-energy orbitals.
The approximate order of filling of atomic orbitals, following the arrows from 1s to 7p. (After 7p the order includes subshells outside the range of the diagram, starting with 8s.) Klechkowski rule 2.svg
The approximate order of filling of atomic orbitals, following the arrows from 1s to 7p. (After 7p the order includes subshells outside the range of the diagram, starting with 8s.)

The principle works very well (for the ground states of the atoms) for the known 118 elements, although it is sometimes slightly wrong. The modern form of the aufbau principle describes an order of orbital energies given by Madelung's rule (or Klechkowski's rule). This rule was first stated by Charles Janet in 1929, rediscovered by Erwin Madelung in 1936, [11] and later given a theoretical justification by V. M. Klechkowski: [13]

  1. Subshells are filled in the order of increasing n + l.
  2. Where two subshells have the same value of n + l, they are filled in order of increasing n.

This gives the following order for filling the orbitals:

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, (8s, 5g, 6f, 7d, 8p, and 9s)

In this list the subshells in parentheses are not occupied in the ground state of the heaviest atom now known (Og, Z = 118).

The aufbau principle can be applied, in a modified form, to the protons and neutrons in the atomic nucleus, as in the shell model of nuclear physics and nuclear chemistry.

Periodic table

Electron configuration table showing blocks. Periodic table blocks spdf (32 column).svg
Electron configuration table showing blocks.

The form of the periodic table is closely related to the atomic electron configuration for each element. For example, all the elements of group 2 (the table's second column) have an electron configuration of [E] ns2 (where [E] is a noble gas configuration), and have notable similarities in their chemical properties. The periodicity of the periodic table in terms of periodic table blocks is due to the number of electrons (2, 6, 10, and 14) needed to fill s, p, d, and f subshells. These blocks appear as the rectangular sections of the periodic table. The single exception is helium, which despite being an s-block atom is conventionally placed with the other noble gasses in the p-block due to its chemical inertness, a consequence of its full outer shell (though there is discussion in the contemporary literature on whether this exception should be retained).

The electrons in the valence (outermost) shell largely determine each element's chemical properties. The similarities in the chemical properties were remarked on more than a century before the idea of electron configuration. [lower-alpha 2]

Shortcomings of the aufbau principle

The aufbau principle rests on a fundamental postulate that the order of orbital energies is fixed, both for a given element and between different elements; in both cases this is only approximately true. It considers atomic orbitals as "boxes" of fixed energy into which can be placed two electrons and no more. However, the energy of an electron "in" an atomic orbital depends on the energies of all the other electrons of the atom (or ion, or molecule, etc.). There are no "one-electron solutions" for systems of more than one electron, only a set of many-electron solutions that cannot be calculated exactly [lower-alpha 3] (although there are mathematical approximations available, such as the Hartree–Fock method).

The fact that the aufbau principle is based on an approximation can be seen from the fact that there is an almost-fixed filling order at all, that, within a given shell, the s-orbital is always filled before the p-orbitals. In a hydrogen-like atom, which only has one electron, the s-orbital and the p-orbitals of the same shell have exactly the same energy, to a very good approximation in the absence of external electromagnetic fields. (However, in a real hydrogen atom, the energy levels are slightly split by the magnetic field of the nucleus, and by the quantum electrodynamic effects of the Lamb shift.)

Ionization of the transition metals

The naïve application of the aufbau principle leads to a well-known paradox (or apparent paradox) in the basic chemistry of the transition metals. Potassium and calcium appear in the periodic table before the transition metals, and have electron configurations [Ar] 4s1 and [Ar] 4s2 respectively, i.e. the 4s-orbital is filled before the 3d-orbital. This is in line with Madelung's rule, as the 4s-orbital has n + l = 4 (n = 4, l = 0) while the 3d-orbital has n + l = 5 (n = 3, l = 2). After calcium, most neutral atoms in the first series of transition metals (scandium through zinc) have configurations with two 4s electrons, but there are two exceptions. Chromium and copper have electron configurations [Ar] 3d5 4s1 and [Ar] 3d10 4s1 respectively, i.e. one electron has passed from the 4s-orbital to a 3d-orbital to generate a half-filled or filled subshell. In this case, the usual explanation is that "half-filled or completely filled subshells are particularly stable arrangements of electrons". However, this is not supported by the facts, as tungsten (W) has a Madelung-following d4 s2 configuration and not d5 s1, and niobium (Nb) has an anomalous d4 s1 configuration that does not give it a half-filled or completely filled subshell. [14]

The apparent paradox arises when electrons are removed from the transition metal atoms to form ions. The first electrons to be ionized come not from the 3d-orbital, as one would expect if it were "higher in energy", but from the 4s-orbital. This interchange of electrons between 4s and 3d is found for all atoms of the first series of transition metals. [lower-alpha 4] The configurations of the neutral atoms (K, Ca, Sc, Ti, V, Cr, ...) usually follow the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, ...; however the successive stages of ionization of a given atom (such as Fe4+, Fe3+, Fe2+, Fe+, Fe) usually follow the order 1s, 2s, 2p, 3s, 3p, 3d, 4s, ...

This phenomenon is only paradoxical if it is assumed that the energy order of atomic orbitals is fixed and unaffected by the nuclear charge or by the presence of electrons in other orbitals. If that were the case, the 3d-orbital would have the same energy as the 3p-orbital, as it does in hydrogen, yet it clearly does not. There is no special reason why the Fe2+ ion should have the same electron configuration as the chromium atom, given that iron has two more protons in its nucleus than chromium, and that the chemistry of the two species is very different. Melrose and Eric Scerri have analyzed the changes of orbital energy with orbital occupations in terms of the two-electron repulsion integrals of the Hartree–Fock method of atomic structure calculation. [15] More recently Scerri has argued that contrary to what is stated in the vast majority of sources including the title of his previous article on the subject, 3d orbitals rather than 4s are in fact preferentially occupied. [16]

In chemical environments, configurations can change even more: Th3+ as a bare ion has a configuration of [Rn] 5f1, yet in most ThIII compounds the thorium atom has a 6d1 configuration instead. [17] [18] Mostly, what is present is rather a superposition of various configurations. [14] For instance, copper metal is poorly described by either an [Ar] 3d10 4s1 or an [Ar] 3d9 4s2 configuration, but is rather well described as a 90% contribution of the first and a 10% contribution of the second. Indeed, visible light is already enough to excite electrons in most transition metals, and they often continuously "flow" through different configurations when that happens (copper and its group are an exception). [19]

Similar ion-like 3dx 4s0 configurations occur in transition metal complexes as described by the simple crystal field theory, even if the metal has oxidation state  0. For example, chromium hexacarbonyl can be described as a chromium atom (not ion) surrounded by six carbon monoxide ligands. The electron configuration of the central chromium atom is described as 3d6 with the six electrons filling the three lower-energy d orbitals between the ligands. The other two d orbitals are at higher energy due to the crystal field of the ligands. This picture is consistent with the experimental fact that the complex is diamagnetic, meaning that it has no unpaired electrons. However, in a more accurate description using molecular orbital theory, the d-like orbitals occupied by the six electrons are no longer identical with the d orbitals of the free atom.

Other exceptions to Madelung's rule

There are several more exceptions to Madelung's rule among the heavier elements, and as atomic number increases it becomes more and more difficult to find simple explanations such as the stability of half-filled subshells. It is possible to predict most of the exceptions by Hartree–Fock calculations, [20] which are an approximate method for taking account of the effect of the other electrons on orbital energies. Qualitatively, for example, the 4d elements have the greatest concentration of Madelung anomalies, because the 4d–5s gap is larger than the 3d–4s and 5d–6s gaps. [21]

For the heavier elements, it is also necessary to take account of the effects of special relativity on the energies of the atomic orbitals, as the inner-shell electrons are moving at speeds approaching the speed of light. In general, these relativistic effects [22] tend to decrease the energy of the s-orbitals in relation to the other atomic orbitals. [23] This is the reason why the 6d elements are predicted to have no Madelung anomalies apart from lawrencium (for which relativistic effects stabilise the p1/2 orbital as well and cause its occupancy in the ground state), as relativity intervenes to make the 7s orbitals lower in energy than the 6d ones.

The table below shows the configurations of the f-block (green) and d-block (blue) atoms. It shows the ground state configuration in terms of orbital occupancy, but it does not show the ground state in terms of the sequence of orbital energies as determined spectroscopically. For example, in the transition metals, the 4s orbital is of a higher energy than the 3d orbitals; and in the lanthanides, the 6s is higher than the 4f and 5d. The ground states can be seen in the Electron configurations of the elements (data page). However this also depends on the charge: a calcium atom has 4s lower in energy than 3d, but a Ca2+ cation has 3d lower in energy than 4s. In practice the configurations predicted by the Madelung rule are at least close to the ground state even in these anomalous cases. [24] The empty f orbitals in lanthanum, actinium, and thorium contribute to chemical bonding, [25] [26] as do the empty p orbitals in transition metals. [27]

Vacant s, d, and f orbitals have been shown explicitly, as is occasionally done, [28] to emphasise the filling order and to clarify that even orbitals unoccupied in the ground state (e.g. lanthanum 4f or palladium 5s) may be occupied and bonding in chemical compounds. (The same is also true for the p-orbitals, which are not explicitly shown because they are only actually occupied for lawrencium in gas-phase ground states.)

Electron shells filled in violation of Madelung's rule [29] (red)
Predictions for elements 109–112 [30]
Period 4 Period 5 Period 6 Period 7
ElementZElectron Configuration ElementZElectron Configuration ElementZElectron Configuration ElementZElectron Configuration
     Lanthanum 57[ Xe ] 6s24f0 5d1  Actinium 89[ Rn ] 7s25f0 6d1
     Cerium 58[ Xe ] 6s24f1 5d1  Thorium 90[ Rn ] 7s25f0 6d2
     Praseodymium 59[ Xe ] 6s2 4f3 5d0  Protactinium 91[ Rn ] 7s25f2 6d1
     Neodymium 60[ Xe ] 6s2 4f4 5d0  Uranium 92[ Rn ] 7s25f3 6d1
     Promethium 61[ Xe ] 6s2 4f5 5d0  Neptunium 93[ Rn ] 7s25f4 6d1
     Samarium 62[ Xe ] 6s2 4f6 5d0  Plutonium 94[ Rn ] 7s2 5f6 6d0
     Europium 63[ Xe ] 6s2 4f7 5d0  Americium 95[ Rn ] 7s2 5f7 6d0
     Gadolinium 64[ Xe ] 6s24f7 5d1  Curium 96[ Rn ] 7s25f7 6d1
     Terbium 65[ Xe ] 6s2 4f9 5d0  Berkelium 97[ Rn ] 7s2 5f9 6d0
     Dysprosium 66[ Xe ] 6s2 4f10 5d0  Californium 98[ Rn ] 7s2 5f10 6d0
     Holmium 67[ Xe ] 6s2 4f11 5d0  Einsteinium 99[ Rn ] 7s2 5f11 6d0
     Erbium 68[ Xe ] 6s2 4f12 5d0  Fermium 100[ Rn ] 7s2 5f12 6d0
     Thulium 69[ Xe ] 6s2 4f13 5d0  Mendelevium 101[ Rn ] 7s2 5f13 6d0
     Ytterbium 70[ Xe ] 6s2 4f14 5d0  Nobelium 102[ Rn ] 7s2 5f14 6d0
Scandium 21[ Ar ] 4s2 3d1  Yttrium 39[ Kr ] 5s2 4d1  Lutetium 71[ Xe ] 6s2 4f14 5d1  Lawrencium 103[ Rn ] 7s2 5f146d0 7p1
Titanium 22[ Ar ] 4s2 3d2  Zirconium 40[ Kr ] 5s2 4d2  Hafnium 72[ Xe ] 6s2 4f14 5d2  Rutherfordium 104[ Rn ] 7s2 5f14 6d2
Vanadium 23[ Ar ] 4s2 3d3  Niobium 41[ Kr ]5s1 4d4  Tantalum 73[ Xe ] 6s2 4f14 5d3  Dubnium 105[ Rn ] 7s2 5f14 6d3
Chromium 24[ Ar ]4s1 3d5  Molybdenum 42[ Kr ]5s1 4d5  Tungsten 74[ Xe ] 6s2 4f14 5d4  Seaborgium 106[ Rn ] 7s2 5f14 6d4
Manganese 25[ Ar ] 4s2 3d5  Technetium 43[ Kr ] 5s2 4d5  Rhenium 75[ Xe ] 6s2 4f14 5d5  Bohrium 107[ Rn ] 7s2 5f14 6d5
Iron 26[ Ar ] 4s2 3d6  Ruthenium 44[ Kr ]5s1 4d7  Osmium 76[ Xe ] 6s2 4f14 5d6  Hassium 108[ Rn ] 7s2 5f14 6d6
Cobalt 27[ Ar ] 4s2 3d7  Rhodium 45[ Kr ]5s1 4d8  Iridium 77[ Xe ] 6s2 4f14 5d7  Meitnerium 109[ Rn ] 7s2 5f14 6d7
Nickel 28[ Ar ] 4s2 3d8 or
[ Ar ]4s1 3d9 (disputed) [31]
  Palladium 46[ Kr ]5s0 4d10  Platinum 78[ Xe ]6s1 4f145d9  Darmstadtium 110[ Rn ] 7s2 5f14 6d8
Copper 29[ Ar ]4s1 3d10  Silver 47[ Kr ]5s1 4d10  Gold 79[ Xe ]6s1 4f145d10  Roentgenium 111[ Rn ] 7s2 5f14 6d9
Zinc 30[ Ar ] 4s2 3d10  Cadmium 48[ Kr ] 5s2 4d10  Mercury 80[ Xe ] 6s2 4f14 5d10  Copernicium 112[ Rn ] 7s2 5f14 6d10

The various anomalies describe the free atoms and do not necessarily predict chemical behavior. Thus for example neodymium typically forms the +3 oxidation state, despite its configuration [Xe] 4f4 5d0 6s2 that if interpreted naïvely would suggest a more stable +2 oxidation state corresponding to losing only the 6s electrons. Contrariwise, uranium as [Rn] 5f3 6d1 7s2 is not very stable in the +3 oxidation state either, preferring +4 and +6. [32]

The electron-shell configuration of elements beyond hassium has not yet been empirically verified, but they are expected to follow Madelung's rule without exceptions until element 120. Element 121 should have the anomalous configuration [Og] 8s25g0 6f0 7d08p1, having a p rather than a g electron. Electron configurations beyond this are tentative and predictions differ between models, [33] but Madelung's rule is expected to break down due to the closeness in energy of the 5g, 6f, 7d, and 8p1/2 orbitals. [30] That said, the filling sequence 8s, 5g, 6f, 7d, 8p is predicted to hold approximately, with perturbations due to the huge spin-orbit splitting of the 8p and 9p shells, and the huge relativistic stabilisation of the 9s shell. [34]

Open and closed shells

In the context of atomic orbitals, an open shell is a valence shell which is not completely filled with electrons or that has not given all of its valence electrons through chemical bonds with other atoms or molecules during a chemical reaction. Conversely a closed shell is obtained with a completely filled valence shell. This configuration is very stable. [35]

For molecules, "open shell" signifies that there are unpaired electrons. In molecular orbital theory, this leads to molecular orbitals that are singly occupied. In computational chemistry implementations of molecular orbital theory, open-shell molecules have to be handled by either the restricted open-shell Hartree–Fock method or the unrestricted Hartree–Fock method. Conversely a closed-shell configuration corresponds to a state where all molecular orbitals are either doubly occupied or empty (a singlet state). [36] Open shell molecules are more difficult to study computationally. [37]

Noble gas configuration

Noble gas configuration is the electron configuration of noble gases. The basis of all chemical reactions is the tendency of chemical elements to acquire stability. Main-group atoms generally obey the octet rule, while transition metals generally obey the 18-electron rule. The noble gases (He, Ne, Ar, Kr, Xe, Rn) are less reactive than other elements because they already have a noble gas configuration. Oganesson is predicted to be more reactive due to relativistic effects for heavy atoms.

1 He 1s2
2 Ne 1s22s2 2p6
3 Ar 1s22s2 2p63s2 3p6
4 Kr 1s22s2 2p63s2 3p64s2 3d10 4p6
5 Xe 1s22s2 2p63s2 3p64s2 3d10 4p65s2 4d10 5p6
6 Rn 1s22s2 2p63s2 3p64s2 3d10 4p65s2 4d10 5p66s2 4f14 5d10 6p6
7 Og 1s22s2 2p63s2 3p64s2 3d10 4p65s2 4d10 5p66s2 4f14 5d10 6p67s2 5f14 6d10 7p6

Every system has the tendency to acquire the state of stability or a state of minimum energy, and so chemical elements take part in chemical reactions to acquire a stable electronic configuration similar to that of its nearest noble gas. An example of this tendency is two hydrogen (H) atoms reacting with one oxygen (O) atom to form water (H2O). Neutral atomic hydrogen has one electron in its valence shell, and on formation of water it acquires a share of a second electron coming from oxygen, so that its configuration is similar to that of its nearest noble gas helium (He) with two electrons in its valence shell. Similarly, neutral atomic oxygen has six electrons in its valence shell, and acquires a share of two electrons from the two hydrogen atoms, so that its configuration is similar to that of its nearest noble gas neon with eight electrons in its valence shell.

Electron configuration in molecules

Electron configuration in molecules is more complex than the electron configuration of atoms, as each molecule has a different orbital structure. The molecular orbitals are labelled according to their symmetry, [lower-alpha 5] rather than the atomic orbital labels used for atoms and monatomic ions; hence, the electron configuration of the dioxygen molecule, O2, is written 1σg2 u2 g2 u2 g2 u4 g2, [38] [39] or equivalently 1σg2 u2 g2 u2 u4 g2 g2. [1] The term 1πg2 represents the two electrons in the two degenerate π*-orbitals (antibonding). From Hund's rules, these electrons have parallel spins in the ground state, and so dioxygen has a net magnetic moment (it is paramagnetic). The explanation of the paramagnetism of dioxygen was a major success for molecular orbital theory.

The electronic configuration of polyatomic molecules can change without absorption or emission of a photon through vibronic couplings.

Electron configuration in solids

In a solid, the electron states become very numerous. They cease to be discrete, and effectively blend into continuous ranges of possible states (an electron band). The notion of electron configuration ceases to be relevant, and yields to band theory.


The most widespread application of electron configurations is in the rationalization of chemical properties, in both inorganic and organic chemistry. In effect, electron configurations, along with some simplified forms of molecular orbital theory, have become the modern equivalent of the valence concept, describing the number and type of chemical bonds that an atom can be expected to form.

This approach is taken further in computational chemistry, which typically attempts to make quantitative estimates of chemical properties. For many years, most such calculations relied upon the "linear combination of atomic orbitals" (LCAO) approximation, using an ever-larger and more complex basis set of atomic orbitals as the starting point. The last step in such a calculation is the assignment of electrons among the molecular orbitals according to the aufbau principle. Not all methods in computational chemistry rely on electron configuration: density functional theory (DFT) is an important example of a method that discards the model.

For atoms or molecules with more than one electron, the motion of electrons are correlated and such a picture is no longer exact. A very large number of electronic configurations are needed to exactly describe any multi-electron system, and no energy can be associated with one single configuration. However, the electronic wave function is usually dominated by a very small number of configurations and therefore the notion of electronic configuration remains essential for multi-electron systems.

A fundamental application of electron configurations is in the interpretation of atomic spectra. In this case, it is necessary to supplement the electron configuration with one or more term symbols, which describe the different energy levels available to an atom. Term symbols can be calculated for any electron configuration, not just the ground-state configuration listed in tables, although not all the energy levels are observed in practice. It is through the analysis of atomic spectra that the ground-state electron configurations of the elements were experimentally determined.

See also


  1. 1 2 In formal terms, the quantum numbers n, l and ml arise from the fact that the solutions to the time-independent Schrödinger equation for hydrogen-like atoms are based on spherical harmonics.
  2. The similarities in chemical properties and the numerical relationship between the atomic weights of calcium, strontium and barium was first noted by Johann Wolfgang Döbereiner in 1817.
  3. Electrons are identical particles, a fact that is sometimes referred to as "indistinguishability of electrons". A one-electron solution to a many-electron system would imply that the electrons could be distinguished from one another, and there is strong experimental evidence that they can't be. The exact solution of a many-electron system is a n-body problem with n  3 (the nucleus counts as one of the "bodies"): such problems have evaded analytical solution since at least the time of Euler.
  4. There are some cases in the second and third series where the electron remains in an s-orbital.
  5. The labels are written in lowercase to indicate that they correspond to one-electron functions. They are numbered consecutively for each symmetry type (irreducible representation in the character table of the point group for the molecule), starting from the orbital of lowest energy for that type.

Related Research Articles

<span class="mw-page-title-main">Atomic orbital</span> Function describing an electron in an atom

In atomic theory and quantum mechanics, an atomic orbital is a function describing the location and wave-like behavior of an electron in an atom. This function can be used to calculate the probability of finding any electron of an atom in any specific region around the atom's nucleus. The term atomic orbital may also refer to the physical region or space where the electron can be calculated to be present, as predicted by the particular mathematical form of the orbital.

<span class="mw-page-title-main">Periodic table</span> Tabular arrangement of the chemical elements ordered by atomic number

The periodic table, also known as the periodic table of the elements, arranges the chemical elements into rows ("periods") and columns ("groups"). It is an icon of chemistry and is widely used in physics and other sciences. It is a depiction of the periodic law, which states that when the elements are arranged in order of their atomic numbers an approximate recurrence of their properties is evident. The table is divided into four roughly rectangular areas called blocks. Elements in the same group tend to show similar chemical characteristics.

In chemistry, a transition metal is a chemical element in the d-block of the periodic table, though the elements of group 12 are sometimes excluded. The lanthanide and actinide elements are called inner transition metals and are sometimes considered to be transition metals as well.

<span class="mw-page-title-main">Atomic radius</span> Measure of the size of an atom

The atomic radius of a chemical element is a measure of the size of its atom, usually the mean or typical distance from the center of the nucleus to the outermost isolated electron. Since the boundary is not a well-defined physical entity, there are various non-equivalent definitions of atomic radius. Four widely used definitions of atomic radius are: Van der Waals radius, ionic radius, metallic radius and covalent radius. Typically, because of the difficulty to isolate atoms in order to measure their radii separately, atomic radius is measured in a chemically bonded state; however theoretical calculations are simpler when considering atoms in isolation. The dependencies on environment, probe, and state lead to a multiplicity of definitions.

<span class="mw-page-title-main">Energy level</span> Different states of quantum systems

A quantum mechanical system or particle that is bound—that is, confined spatially—can only take on certain discrete values of energy, called energy levels. This contrasts with classical particles, which can have any amount of energy. The term is commonly used for the energy levels of the electrons in atoms, ions, or molecules, which are bound by the electric field of the nucleus, but can also refer to energy levels of nuclei or vibrational or rotational energy levels in molecules. The energy spectrum of a system with such discrete energy levels is said to be quantized.

<span class="mw-page-title-main">Ionization energy</span> Energy needed to remove an electron

In physics and chemistry, ionization energy (IE) (American English spelling), ionisation energy (British English spelling) is the minimum energy required to remove the most loosely bound electron of an isolated gaseous atom, positive ion, or molecule. The first ionization energy is quantitatively expressed as

An extended periodic table theorizes about chemical elements beyond those currently known and proven. The element with the highest atomic number known is oganesson (Z = 118), which completes the seventh period (row) in the periodic table. All elements in the eighth period and beyond thus remain purely hypothetical.

A period 4 element is one of the chemical elements in the fourth row of the periodic table of the chemical elements. The periodic table is laid out in rows to illustrate recurring (periodic) trends in the chemical behaviour of the elements as their atomic number increases: a new row is begun when chemical behaviour begins to repeat, meaning that elements with similar behaviour fall into the same vertical columns. The fourth period contains 18 elements beginning with potassium and ending with krypton – one element for each of the eighteen groups. It sees the first appearance of d-block in the table.

<span class="mw-page-title-main">Octet rule</span> Chemical rule of thumb

The octet rule is a chemical rule of thumb that reflects the theory that main-group elements tend to bond in such a way that each atom has eight electrons in its valence shell, giving it the same electronic configuration as a noble gas. The rule is especially applicable to carbon, nitrogen, oxygen, and the halogens; although more generally the rule is applicable for the s-block and p-block of the periodic table. Other rules exist for other elements, such as the duplet rule for hydrogen and helium, and the 18-electron rule for transition metals.

Hund's rule of maximum multiplicity is a rule based on observation of atomic spectra, which is used to predict the ground state of an atom or molecule with one or more open electronic shells. The rule states that for a given electron configuration, the lowest energy term is the one with the greatest value of spin multiplicity. This implies that if two or more orbitals of equal energy are available, electrons will occupy them singly before filling them in pairs. The rule, discovered by Friedrich Hund in 1925, is of important use in atomic chemistry, spectroscopy, and quantum chemistry, and is often abbreviated to Hund's rule, ignoring Hund's other two rules.

<span class="mw-page-title-main">Valence electron</span> An electron in the outer shell of an atoms energy levels

In chemistry and physics, valence electrons are electrons in the outermost shell of an atom, and that can participate in the formation of a chemical bond if the outermost shell is not closed. In a single covalent bond, a shared pair forms with both atoms in the bond each contributing one valence electron.

In chemistry, orbital hybridisation is the concept of mixing atomic orbitals to form new hybrid orbitals suitable for the pairing of electrons to form chemical bonds in valence bond theory. For example, in a carbon atom which forms four single bonds the valence-shell s orbital combines with three valence-shell p orbitals to form four equivalent sp3 mixtures in a tetrahedral arrangement around the carbon to bond to four different atoms. Hybrid orbitals are useful in the explanation of molecular geometry and atomic bonding properties and are symmetrically disposed in space. Usually hybrid orbitals are formed by mixing atomic orbitals of comparable energies.

In atomic physics, a term symbol is an abbreviated description of the total spin and orbital angular momentum quantum numbers of the electrons in a multi-electron atom. So while the word symbol suggests otherwise, it represents an actual value of a physical quantity.

<span class="mw-page-title-main">Aufbau principle</span> Principle of atomic physics

The Aufbau principle, also called the Aufbau rule, states that in the ground state of an atom or ion, electrons first fill subshells of the lowest available energy, then fill subshells of higher energy. For example, the 1s subshell is filled before the 2s subshell is occupied. In this way, the electrons of an atom or ion form the most stable electron configuration possible. An example is the configuration 1s2 2s2 2p6 3s2 3p3 for the phosphorus atom, meaning that the 1s subshell has 2 electrons, and so on.

<span class="mw-page-title-main">Hund's rules</span> Rules to determine the ground state of an atom

In atomic physics and quantum chemistry, Hund's rules refers to a set of rules that German physicist Friedrich Hund formulated around 1925, which are used to determine the term symbol that corresponds to the ground state of a multi-electron atom. The first rule is especially important in chemistry, where it is often referred to simply as Hund's Rule.

This page shows the electron configurations of the neutral gaseous atoms in their ground states. For each atom the subshells are given first in concise form, then with all subshells written out, followed by the number of electrons per shell. For phosphorus as an example, the concise form is [Ne] 3s2 3p3. Here [Ne] refers to the core electrons which are the same as for the element neon (Ne), the last noble gas before phosphorus in the periodic table. The valence electrons are written explicitly for all atoms.

Core electrons are the electrons in an atom that are not valence electrons and do not participate in chemical bonding. The nucleus and the core electrons of an atom form the atomic core. Core electrons are tightly bound to the nucleus. Therefore, unlike valence electrons, core electrons play a secondary role in chemical bonding and reactions by screening the positive charge of the atomic nucleus from the valence electrons.

In chemistry and atomic physics, an electron shell may be thought of as an orbit that electrons follow around an atom's nucleus. The closest shell to the nucleus is called the "1 shell" (also called the "K shell"), followed by the "2 shell" (or "L shell"), then the "3 shell" (or "M shell"), and so on farther and farther from the nucleus. The shells correspond to the principal quantum numbers (n = 1, 2, 3, 4 ...) or are labeled alphabetically with the letters used in X-ray notation (K, L, M, ...). A useful guide when understanding electron shells in atoms is to note that each row on the conventional periodic table of elements represents an electron shell.

The d electron count or number of d electrons is a chemistry formalism used to describe the electron configuration of the valence electrons of a transition metal center in a coordination complex. The d electron count is an effective way to understand the geometry and reactivity of transition metal complexes. The formalism has been incorporated into the two major models used to describe coordination complexes; crystal field theory and ligand field theory, which is a more advanced version based on molecular orbital theory. However the d electron count of an atom in a complex is often different from the d electron count of a free atom or a free ion of the same element.


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