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How do I read an electron configuration table?Are you
and need to know how to place the
around the ? If so, you will need to know how to read an element's electron configuration table. Follow these easy directions to learn how!What is an electron configuration table?An electron configuration table is a type of code that describes how many electrons are in each energy level of an atom and how the electrons are arranged within each energy level. It packs a lot of information into a little space and it takes a little practice to read. For example, this is the electron configuration table for :What do all those numbers and letters mean?Each row of an electron configuration table is sort of like a sentence. Each 'sentence' is made up of smaller 'words'. Each 'word' follows this format:The first number is the energy level. We can tell right away that an atom of gold contains 6 energy levels.The lowercase letter is the sub-shell. The sub-shells are named s, p, d and f. The number of available sub-shells increases as the energy level increases. For example, the first energy level only contains an s&sub-shell while the second energy level contains both an s&sub-shell and a p&sub-shell.The number in superscript is the number of electrons in a sub-shell. Each sub-shell can hold only a certain number of electrons. The s&sub-shell can hold no more than 2 electrons, the p&sub-shell can hold 6, the d&sub-shell can hold 10 and the f&sub-shell can hold as many as 14.How can I use the electron configuration table to tell me...How many energy levels does an atom have?Since the electron configuration table lists each energy level by row, you can tell how many energy levels there are by seeing how many rows there are. As was mentioned earlier, an atom of gold contains six energy levels, as shown below:How many electrons are in each energy level?The total number of electrons in an energy level is the sum of the electrons in each sub-shell of that energy level. Just add the numbers in superscript together to find the number of electrons in an energy level. The number of electrons in each energy level in an atom of gold is shown below:How many electrons are in an atom's outer energy level?This is just a combination of the previous two examples. Use the electron configuration to find that atom's highest energy level and then add up the numbers in superscript to find the number of electrons that are in it. There is one electron in the outer energy level of an atom of gold, as shown below:
Related Pages:From Wikipedia, the free encyclopedia
and , the electron configuration is the distribution of
(or other physical structure) in
or . For example, the electron configuration of the
atom is 1s2 2s2 2p6.
Electronic configurations describe electrons as each moving independently in an orbital, in an average field created by all other orbitals. Mathematically, configurations are described by
According to the laws of , for systems with only one electron, an energy is associated with each electron configuration and, upon certain conditions, electrons are able to move from one configuration to another by the emission or absorption of a
of energy, in the form of a .
Knowledge of the electron configuration of different atoms is useful in understanding the structure of the
of elements. The concept is also useful for describing the chemical bonds that hold atoms together. In bulk materials, this same idea helps explain the peculiar properties of
Electron configuration was first conceived of under the
of the atom, and it is still common to speak of shells and subshells despite the advances in understanding of the
nature of electrons.
An electron shell is the set of
that share the same , n (the number before the letter in the orbital label), that electrons may occupy. An atom's nth electron shell can accommodate 2n2 electrons, e.g. the first shell can accommodate 2 electrons, the second shell 8 electrons, and the third shell 18 electrons. The factor of two arises because the allowed states are doubled due to —each
admits up to two otherwise identical electrons with opposite spin, one with a spin +1/2 (usually noted by an up-arrow) and one with a spin -1/2 (with a down-arrow).
A subshell is the set of states defined by a common , l, within a shell. The values l = 0, 1, 2, 3 correspond to the s, p, d, and f labels, respectively. 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 arises from the equations of quantum mechanics, in particular the , which states that no two electrons in the same atom can have the same values of the four .
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 orbital labels (e.g. for
the sequence 1s, 2s, 2p, 3s, 3p) with the number of electrons assigned to each orbital (or set of orbitals sharing the same label) placed as a superscript. For example,
has one electron in the s-orbital of the first shell, so its configuration is written 1s1.
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").
( 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, since all but the last few subshells are identical to those of one or another of the . Phosphorus, for instance, differs from
(1s2 2s2 2p6) only by the presence of a third shell. Thus, the electron configuration of neon is pulled out, and phosphorus is written as follows: [Ne] 3s2 3p3. 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
ground state can be written as either [Ar] 4s2 3d2 or [Ar] 3d2 4s2. The first notation follows the order based on the
for the configurati 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 t 3d is filled before 4s in the sequence Ti4+, Ti3+, Ti2+, Ti+, Ti.
The superscript 1 for a singly occupied orbital is not compulsory. 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
(IUPAC) recommends a normal typeface (as used here). The choice of letters originates from a now-obsolete system of categorizing
as "sharp", "principal", "diffuse" and "fundamental" (or "fine"), based on their observed : their modern usage indicates orbitals with an , 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.
The electron configurations of molecules are written in a similar way, except that
labels are used instead of atomic orbital labels (see below).
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 . Any other configuration is an .
As an example, the ground state configuration of the
atom is 1s22s22p63s, as deduced from the Aufbau principle (see below). The first excited state is obtained by promoting a 3s electron to the 3p orbital, to obtain the 1s22s22p63p configuration, abbreviated as the 3p level. Atoms can move from one configuration to another by absorbing or emitting energy. In a
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
(such as 3s for sodium) involves energies corresponding to
of visible or
light. The excitation of
is possible, but requires much higher energies, generally corresponding to
photons. This would be the case for example to excite a 2p electron to the 3s level and form the excited 1s22s22p53s2 configuration.
The remainder of this article deals only with the ground-state configuration, often referred to as "the" configuration of an atom or molecule.
(1923) was the first to propose that the
in the properties of the elements might be explained by the electronic structure of the atom. His proposals were based on the then current
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:
was given as 2.4.4.6 instead of 1s2 2s2 2p6 3s2 3p4 (2.8.6).
The following year,
incorporated
third quantum number into the description of electron shells, and correctly predicted the shell structure of sulfur to be 2.8.6. However neither Bohr's system nor Stoner's could correctly describe the changes in
Bohr was well aware of this shortcoming (and others), and had written to his friend
to ask for his help in saving quantum theory (the system now known as ""). Pauli realized that the Zeeman effect must be due only to the outermost electrons of the atom, and 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
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 , published in 1926, gave three of the four quantum numbers as a direct consequence of its solution for the hydrogen atom: 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), see below) for the order in which atomic orbitals are filled with electrons.
Aufbau, "building up, construction") was an important part of Bohr's original concept of electron configuration. It may be stated as:
a maximum of two electrons are put into orbitals in the order of increasing orbital energy: the lowest-energy orbitals 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 orbitals outside the range of the diagram, starting with 8s.)
The principle works very well (for the ground states of the atoms) for the first 18 elements, then decreasingly well for the following 100 elements. 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
in 1929, rediscovered by
in 1936, and later given a theoretical justification by
Orbitals are filled in the order of increasing n+l;
Where two orbitals 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 orbitals in parentheses are not occupied in the ground state of the heaviest atom now known (, Z = 118).
The Aufbau principle can be applied, in a modified form, to the
in the , as in the
Electron configuration table
The form of the
is closely related to the electron configuration of the atoms of the elements. For example, all the elements of
have an electron configuration of [E] ns2 (where [E] is an
configuration), and have notable similarities in their chemical properties. In general, the periodicity of the periodic table in terms of
is clearly due to the number of electrons (2, 6, 10, 14...) needed to fill s, p, d, and f subshells.
The outermost electron shell is often referred to as the "valence shell" and (to a first approximation) determines the chemical properties. It should be remembered that the similarities in the chemical properties were remarked on more than a century before the idea of electron configuration. It is not clear how far Madelung's rule explains (rather than simply describes) the periodic table, although some properties (such as the common +2
in the first row of the transition metals) would obviously be different with a different order of orbital filling.
The Aufbau principle rests on a fundamental postulate that the order of orbital energies is fixed, both for a given element and betwee 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 (although there are mathematical approximations available, such as the ).
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 , 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
effects of the .)
The na?ve application of the Aufbau principle leads to a well-known
(or apparent paradox) in the basic chemistry of the .
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 (Sc-Zn) have configurations with two 4s electrons, but there are two exceptions.
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".
The apparent paradox arises when electrons are removed from the transition metal atoms to form . 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. 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 doesn't. There is no special reason why the Fe2+ ion should have the same electron configuration as the chromium atom, given that
has two more protons in its nucleus than chromium, and that the chemistry of the two species is very different. Melrose and
have analyzed the changes of orbital energy with orbital occupations in terms of the two-electron repulsion integrals of the
of atomic structure calculation.
Similar ion-like 3dx4s0 configurations occur in
as described by the simple , even if the metal has  0. For example,
can be described as a chromium atom (not ion) surrounded by six
. 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 , meaning that it has no unpaired electrons. However, in a more accurate description using , the d-like orbitals occupied by the six electrons are no longer identical with the d orbitals of the free atom.
There are several more exceptions to
among the heavier elements, and it is more and more difficult to resort to simple explanations, such as the stability of half-filled subshells. It is possible to predict most of the exceptions by Hartree–Fock calculations, which are an approximate method for taking account of the effect of the other electrons on orbital energies. For the heavier elements, it is also necessary to take account of the
on the energies of the atomic orbitals, as the inner-shell electrons are moving at speeds approaching the . In general, these relativistic effects tend to decrease the energy of the s-orbitals in relation to the other atomic orbitals. The table below 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 3 and in the lanthanides, the 6s is higher than the 4f and 5d. The ground states can be seen in the .
Electron shells filled in violation of Madelung's rule (red)
Electron Configuration
Electron Configuration
Electron Configuration
Electron Configuration
[] 6s2 5d1
[] 7s2 6d1
[] 6s2 4f1 5d1
[] 7s2 6d2
[] 6s2 4f3
[] 7s2 5f2 6d1
[] 6s2 4f4
[] 7s2 5f3 6d1
[] 6s2 4f5
[] 7s2 5f4 6d1
[] 6s2 4f6
[] 7s2 5f6
[] 6s2 4f7
[] 7s2 5f7
[] 6s2 4f7 5d1
[] 7s2 5f7 6d1
[] 6s2 4f9
[] 7s2 5f9
[] 4s2 3d1
[] 5s2 4d1
[] 6s2 4f14 5d1
[] 7s2 5f14 7p1
[] 4s2 3d2
[] 5s2 4d2
[] 6s2 4f14 5d2
[] 7s2 5f14 6d2
[] 4s2 3d3
[] 5s1 4d4
[] 6s2 4f14 5d3
[] 4s1 3d5
[] 5s1 4d5
[] 6s2 4f14 5d4
[] 4s2 3d5
[] 5s2 4d5
[] 6s2 4f14 5d5
[] 4s2 3d6
[] 5s1 4d7
[] 6s2 4f14 5d6
[] 4s2 3d7
[] 5s1 4d8
[] 6s2 4f14 5d7
[] 4s2 3d8 or
[] 4s1 3d9 ()
[] 6s1 4f14 5d9
[] 4s1 3d10
[] 5s1 4d10
[] 6s1 4f14 5d10
[] 4s2 3d10
[] 5s2 4d10
[] 6s2 4f14 5d10
The electron-shell configuration of elements beyond
has not yet been empirically verified, but they are expected to follow Madelung's rule without exceptions until .
In , the situation becomes more complex, as each molecule has a different orbital structure. The
are labelled according to their , rather than the
labels used for atoms and monatomic ions: hence, the electron configuration of the
molecule, O2, is written 1σg2 1σu2 2σg2 2σu2 3σg2 1πu4 1πg2, or equivalently 1σg2 1σu2 2σg2 2σu2 1πu4 3σg2 1πg2. The term 1πg2 represents the two electrons in the two degenerate π*-orbitals (antibonding). From , these electrons have parallel spins in the , and so dioxygen has a net
(it is ). The explanation of the paramagnetism of dioxygen was a major success for .
The electronic configuration of polyatomic molecules can change without absorption or emission of a photon through .
In a , the electron states become very numerous. They cease to be discrete, and effectively blend into continuous ranges of possible states (an ). The notion of electron configuration ceases to be relevant, and yields to .
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 form of , have become the modern equivalent of the
concept, describing the number and type of chemical bonds that an atom can be expected to form.
This approach is taken further in , which typically attempts to make quantitative estimates of chemical properties. For many years, most such calculations relied upon the "" (LCAO) approximation, using an ever larger and more complex
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 calculational chemistry rely on electron configuration:
(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
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 . In this case, it is necessary to supplement the electron configuration with one or more , 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.
Wikimedia Commons has media related to .
Discusses the limits of the periodic table
, , 2nd ed. (the "Gold Book") (1997). Online corrected version:  (2006–) "".
In formal terms, the
n, l and ml arise from the fact that the solutions to the time-independent
are based on .
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Electrons are , 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
with n ≥ 3 (the nucleus counts as one of the "bodies"): such problems have evaded
since at least the time of .
There are some cases in the second and third series where the electron remains in an s-orbital.
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The labels are written in lowercase to indicate that they correspond to one-electron functions. They are numbered consecutively for each symmetry type ( in the
for the molecule), starting from the orbital of lowest energy for that type.
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