Electronic Configuration

In Physical of the particles as in Chemistry, the electronic configuration is the arrangement of the electron S in a Atome, a Molécule or another body. Precisely, it is the position of the electrons in orbital a atomic, molecular or other electronic orbital forms of .

Electronic configuration of an atom

the following discussion requires a minimal knowledge of the atomic article Orbitale.

Summary of the quantum numbers

The state of an electron in an atom is provided by four quantum numbers. The three first provide the properties of the atomic Orbitale in which the electron in question is (a more detailed explanation is given in this last article).

the principal quantum Number, noted N , can take any higher whole value or equalizes to 1. It partly represents total energy of orbital, and by extension its distance general compared to the Atomic nucleus;
the azimuth quantum Number, noted L , can take any whole value in the interval $0 \ the L \ the n-1$ . It determines the angular momentum of the orbital one;
the magnetic quantum Number, noted m , can take any whole value included/understood in the interval $-l \ the m \ the l$ . This number determines the energy change of a atomic Orbitale due to an external magnetic field (Effet Zeeman);
the quantum number of spin, noted S , can take only values +1/2 or -1/2 (sometimes simply called in top and in bottom ). The spin is an intrinsic property of the electron and is independent of the other numbers. The spin (in conjunction with the angular momentum) partly determines the magnetic moment of the dipole of the electron.

Layers and underlayers

The layers and underlayers are identified like the above-mentioned quantum numbers, and NOT by the distance from the electrons to the core. In fact, in the atoms of big size, the roadbases than the second layer are superimposed (see the paragraph principle Afbau).

The electronic states with the same value of N are in relation, and it is said that they belong to same the electronic Couche . The electronic states with the same value of N and also of L belong to same the electronic Sous-couche . If the states also share the same value of m , it is said that they belong to same the atomic Orbitale . And because the electrons have only two values possible of spin (quantum number S ), orbital atomic cannot contain more than two electrons (Principe of exclusion of Pauli).

On the basis of these definitions, one from of deduced full capacity in electrons for

a layer: 2 N ²
an underlayer: 4 L +2
a quantum box: 2

Example

Here electronic configuration of a fifth layer filled:

This information can be noted as follows: 5s² 5p⁶ 5d¹⁰ 5f¹⁴ 5g¹⁸ (see below for more details on the notation).

The names of the underlayers (or orbital) S, p, D and F come from a system of categorization starting from the spectral lines known as (in English) sharp , principal , diffuses or fundamental , based on the observation of their fine Structure. When the first four types of orbital were described, they were associated with these four types of spectral lines. Designation G comes from the continuation according to the alphabetical order. The layers with more than 6 underlayers are theoretically possible, but for the moment no chemical element has electrons in an underlayer G . The first element which would have an electron in an underlayer G would be the Unbiunium (Atomic number 121)

Notation

The physicists and the chemists use a standard notation to describe the configurations of the electrons of the atom. According to this notation, an underlayer is noted in the form nx^e, where N is the number of the layer, x is the name of the underlayer and E is the number of electrons in the underlayer. The underlayers of the atom are written in the order ascending of energy - in other words, the sequence according to which they are filled (see the principle of Aufbau below).

For example, the Hydrogène in a fundamental state has an electron in the underlayer S of the first layer, therefore its configuration is noted 1s¹. The Lithium has two electrons in the underlayer 1s and one under 2s layer (of higher energy than the preceding one). Its configuration is thus noted 1s² 2s¹. For the Phosphorus (atomic number 15), that gives: 1s² 2s² 2p⁶ 3s² 3p³.

For the atoms having several electronics layers completely filled, the notation can become very long. One can then shorten the notation by indicating that the configuration of the electrons of heart (electrons of an electron shell completely filled, except for the layers D or F) have a configuration identical to that of the first rare gas which precedes the element. Phosphorus, for example, does not differ from the Néon (1s² 2s² 2p⁶) only by the presence of a third layer. Thus the electronic configuration of neon is withdrawn, and phosphorus is noted as follows: 3s² 3p³.

A manner even simpler to note the number of electrons in each layer is for example (always for phosphorus): 2-8-5

Aufbau principle

To state fundamental of atom (the state in which it is usually), the electronic configuration follows the Principe Aufbau. According to this principle, the electrons place in electronic states of increasing energy . For example, the first electron is placed in the state presenting the energy level low, the second electron in the following state of lower energy, etc

The order in which the underlayers are filled is defined according to the rule of Klechkowski, represented on the following figure, the diagram of Klechkowski:

(There must be also a 5g over the 6f)

A pair of electrons with identical spins is slightly less less expensive in energy than a pair of electrons with opposite spins. Or to say it more precisely, to pair two electrons of spin opposed on the very orbital one (since two electrons on the same orbital one must have opposite spins) more energy costs than to put first electrons on another orbital. That made that the electrons " préfèrent" to occupy first of all of orbital different from/to each other. This preference (which corresponds to a more stable configuration or less energetics) appears itself if an underlayer with $l>0$ (i.e. who contains more than the one orbital one) is not complete. For example, if an underlayer p contains four electrons, two electrons will occupy orbital, but the two others will occupy each one one of orbital remaining, and their spin will be the same one. This phenomenon is described by the Règle of Hund.

The Aufbau principle can be applied, in a modified form, with the Proton S and the Neutron S in the Atomic nucleus.

Exceptions

An underlayer D which is filled with half (thus which has 5 electrons out of 10 possible) is more stable than the underlayer S of the next layer. That is possible because it costs less energy to maintain an electron in an underlayer D with half filled than in an underlayer S filled. For example, the Cuivre (atomic number 29) has a configuration of 3d¹⁰ 4s¹, and not 3d⁹ 4s² as one could suppose it according to the Aufbau principle. Same manner, the Chrome (atomic number 24) has a configuration of 3d⁵ 4s¹, and not 3d⁴ 4s².

That can be included/understood more easily by traversing the illustrated electronic configuration with.

Correspondence with the structure of the periodic table

The electronic configuration is closely related to the structure of the periodic Table. The chemical properties of an atom are largely determined by the arrangement of the electrons in la' layer externe' (or layer of valence ) (although other factors, like the atomic Rayon, the Atomic mass and an increased accessibility in other electronic states also contribute to the chemistry of the elements when the size of the atom increases).

Configuration of the electrons in the molecules

In the molecules, the situation becomes more complex, because each molecule has a different orbital structure. One will see the article on the molecular Orbitale and the method of linear Combinaison of orbital atomic for an introduction on this subject, and the article on the computational Chimie for a thorough knowledge.

Electronic configuration in the solids

In a solid, the electronic states become very numerous. They cease being discrete, and mix together in a continuous extent of possible states (a electronic Bande). The concept of configuration electronic ceases being relevant, and leaves the place to the Théorie of the bands.

See too

The electronic configuration of a chemical element describes the distribution of its electron S between different the orbital, i.e. the layers where the electrons are moving.

The electronic configuration affects the behavior of the Atome in the chemical reactions.

The electrons are laid out around the core so that their energy level is lowest, by respecting the Principe of exclusion of Pauli, the Règle of Klechkowsky as well as the Règle of Hund.

With final, one can write the electronic structure of each atom in his fundamental state. For that, one writes the symbol of the element followed by a series comprising the underlayers of each energy level (orbital) and the number of electrons which they have while exposing.

Example of the Carbon (Z=6): ₆C: 1s²2s²2p²

Example of the Iron (Z=26): ₂₆Fe: 1s²2s²2p⁶3s²3p⁶3d⁶4s² -->

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