Name: octet rule exceptions in lewis structures
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Exceptions to the Octet Rule

La règle d'octet explique la liaison chimique dans les composés du groupe principal en prédisant que chaque atome atteint une configuration à 8 électrons. Cependant, il y a trois exceptions majeures à cette règle. La première exception est les espèces d'électrons impairs.La plupart des molécules et des ions ont un nombre pair d'électrons. Cependant, certaines molécules appelées radicaux, ont un ou plusieurs électrons non appariés. Les radicaux ayant un nombre impair d'électrons non appariés ne peuvent pas atteindre les octets.L'anion superoxyde, un radical avec un électron non apparié, a 13 électrons de valence. Il peut être représenté par deux structures contributives où un oxygène n'a que sept électrons et ne peut donc pas atteindre un octet. La deuxième exception est les atomes qui forment un octet incomplet.Par exemple, l'hydrogène, l'hélium et le lithium ont tendance à atteindre un duo, alors que les éléments du groupe 2 et 13 comme le béryllium et le bore forment souvent des molécules avec quatre et six électrons autour d'eux, respectivement. Considérons le chlorure d'aluminium, qui a 24 électrons de valence. Alors que tous les atomes de chlore atteignent l'octet, l'aluminium n'obtient que 6 électrons de valence un octet incomplet.Bien que le chlorure d'aluminium soit stable, il réagit avec des molécules comme l'ammoniac qui ont une paire d'électrons non partagée. L'azote dans l'ammoniac donne sa seule paire à l'aluminium, formant une liaison spéciale appelée covalente coordonnée, ou datif, liaison. La troisième exception concerne les éléments pouvant accueillir plus de 8 électrons de valence, ou un octet étendu.Ces éléments sont situés dans la troisième rangée du tableau périodique et ci-dessous. Les éléments, tels que le phosphore, le soufre ou l'iode, ont accès aux orbitales d, ce qui leur permet d'accueillir plus de 8 électrons de valence souvent jusqu'à 12 ou 14. Considérons l'anion tétrachloroiodure, qui a 36 électrons de valence.Même après avoir assigné les paires d'électrons de liaison et satisfait l'octet pour tous les atomes, 4 électrons de valence restent non affectés. Ces électrons sont placés sur l'atome d'iode central, ce qui donne un octet élargi avec 12 électrons. Les molécules avec plus de 8 électrons de valence autour de l'atome central sont appelées hypervalentes.Rappelez-vous, les éléments de la deuxième rangée du tableau périodique, tels que le carbone ou l'oxygène, n'ont que des orbitales s et p et ne forment jamais de composés hypervalents, car collectivement, ils ne peuvent contenir que 8 électrons de valence.

exceptions à la règle de l'octet

Beaucoup de molécules covalentes ont des atomes centraux qui n'ont pas huit électrons dans leurs structures de Lewis. Ces molécules se divisent en trois catégories : 
 
 <ol>
 <li>Les molécules d'électrons impairs ont un nombre impair d'électrons de valence et elles ont donc un électron non apparié. </li>
 <li>Les molécules avec un déficit en électrons ont un atome central avec moins d'électrons que nécessaire pour une configuration de gaz noble. </li>
 <li>Les molécules hypervalentes ont un atome central qui a plus d'électrons que nécessaire pour une configuration de gaz noble. </li>
 </ol>
 
 <h4>Molécules d'électrons impairs </h4>
 
 Les molécules qui contiennent un nombre impair d'électrons sont appelées des radicaux. Le monoxyde d'azote, NO, est un exemple de molécule d'électrons impairs ; il est produit dans les moteurs à combustion interne lorsque l'oxygène et l'azote réagissent à des températures élevées. 
 
 Pour tracer la structure de Lewis correspondant à une molécule d'électrons impairs comme NO, les étapes suivantes sont prises en compte : 
 
 <ol>
 <li>Déterminez le nombre total d'électrons de valence (couche externe). La somme des électrons de valence est de 5 (de N) + 6 (de O) = 11. Le nombre impair indique qu'il s'agit d'un radical libre, où chaque atome n'a pas huit électrons sur sa couche de valence. </li>
 <li>Dessinez la structure du squelette de la molécule. Une structure du squelette avec une liaison simple N–O peut facilement être tracée. </li>
 <li>Répartissez les électrons restants sous forme de doublets libres sur les atomes terminaux. Dans ce cas, il n'y a pas d'atome central, donc les électrons sont répartis autour des deux atomes. Huit électrons sont attribués à l'atome le plus électronégatif dans ces situations ; ainsi, l'oxygène a la couche de valence remplie : 
 <img alt="Figure1" src="https://www.jovecore.com/CDNSource/lessons/1066/1066Image_1.svg" style="height: 40px; display: block; margin: 0 auto;" /></li>
 <li>Placez tous les électrons restants sur l'atome central. Puisqu'il n'y a plus d'électrons restants, cette étape ne s'applique pas. </li>
 <li>Réorganisez les électrons pour créer des liaisons multiples avec l'atome central afin d'obtenir des octets dans la mesure du possible. Bien qu'une molécule d'électrons impairs ne puisse pas avoir un octet pour chaque atome, chaque atome devrait avoir des électrons aussi proches que possible d'un octet. Dans ce cas, l'azote n'a que cinq électrons autour de lui. Pour se rapprocher d'un octet dans le cas de l'azote, un des doublets libres de l'oxygène est utilisé pour former une double liaison NO. (Un autre doublet libre d'électrons ne peut pas être prélevé à l'oxygène pour former une liaison triple car l'azote aurait alors neuf électrons :) 
 <img alt="Figure2" src="https://www.jovecore.com/CDNSource/lessons/1066/1066Image_2.svg" style="height: 40px; display: block; margin: 0 auto;" /></li>
 </ol>
 
 <h4>Molécules avec un déficit en électrons </h4>
 
 Certaines molécules contiennent cependant des atomes centraux qui n'ont pas de couche de valence remplie. En général, il s'agit de molécules avec des atomes centraux des groupes 2 et 13, les atomes externes que sont l'hydrogène ou enfin d'autres atomes qui ne forment pas de liaisons multiples. Par exemple, dans les structures de Lewis du dihydrure de béryllium, BeH2, et du trifluorure de bore, BF3, les atomes de béryllium et de bore n'ont chacun que quatre et six électrons, respectivement. Il est possible de tracer une structure avec une liaison double entre un atome de bore et un atome de fluor dans BF3, ce qui satisfait à la règle de l'octet, mais des preuves expérimentales indiquent que les longueurs de la liaison sont plus proches de celles auxquelles on s'attend pour des liaisons simples B–F. Cela suggère que la meilleure structure de Lewis possède trois liaisons simples B–F et un bore avec un déficit en électrons. La réactivité du composé est également compatible avec un bore ayant un déficit en électrons. Cependant, les liaisons B–F sont légèrement plus courtes que ce qui est en fait prévu pour des liaisons simples B–F, ce qui indique qu'un peu du caractère d'une double liaison se trouve dans la molécule réelle. 
 
 <img alt="Figure3" src="https://www.jovecore.com/CDNSource/lessons/1066/1066Image_3.svg" style="height: 80px;" />
 
 Un atome comme l'atome de bore dans BF3, qui n'a pas huit électrons, est très réactif. Il se combine facilement avec une molécule contenant un atome ayant un doublet libre d'électrons. Par exemple, NH3 réagit avec BF3 parce que le doublet libre sur l'azote peut être partagé avec l'atome de bore : 
 
 <img alt="Figure4" src="https://www.jovecore.com/CDNSource/lessons/1066/1066Image_4.svg" style="height: 80px;" />
 
 <h4>Molécules hypervalentes </h4>
 
 Les éléments de la deuxième période du tableau périodique (n = 2) ne peuvent accueillir que huit électrons dans les orbitales de leur couche de valence car ils n'ont que quatre orbitales de valence (une orbitale 2s et trois orbitales 2p). Les éléments de la troisième période et des périodes supérieures (n ≥ 3) ont plus de quatre orbitales de valence et peuvent partager plus de quatre doublets d'électrons avec d'autres atomes parce qu'ils ont des orbitales d vides dans la même couche. Les molécules formées à partir de ces éléments sont parfois appelées des molécules hypervalentes, telles que PCl5 et SF6. Dans PCl5, l'atome central, le phosphore, partage cinq doublets d'électrons. Dans SF6, le soufre partage six doublets d'électrons. 
 
 <img alt="Figure5" src="https://www.jovecore.com/CDNSource/lessons/1066/1066Image_5.svg" style="height: 110px;" />
 
 Dans certaines molécules hypervalentes, comme IF5 et XeF4, certains électrons sur la couche externe de l'atome central sont des doublets libres : 
 
 <img alt="Figure6" src="https://www.jovecore.com/CDNSource/lessons/1066/1066image_6.svg" style="height: 110px;" />
 
 Dans les structures de Lewis pour ces molécules, il reste des électrons après que les couches de valence des atomes externes aient été remplies avec huit électrons. Ces électrons supplémentaires doivent être attribués à l'atome central. 
 
 Ce texte est adapté de <a href="https://openstax.org/books/chemistry-2e/pages/7-3-lewis-symbols-and-structures">Openstax, Chimie 2e, Section 7.3 : Symboles et structures de Lewis.</a>

Exceptions à la règle de l'octet

De nombreuses molécules covalentes ont des atomes centraux qui n’ont pas huit électrons dans leurs structures de Lewis. Ces molécules se répartissent en trois catégories :

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octet rule exceptions in lewis structures

Some covalent molecules do not follow the octet rule in their Lewis structures. Their central atoms can fall into three groups: odd-electron molecules, electron-deficient molecules, and hypervalent molecules. These cases show why a Lewis structure may not always give every atom eight valence electrons.

Odd-electron molecules have an odd number of valence electrons, so one electron remains unpaired. Molecules with an odd number of electrons are called radicals. Nitric oxide, NO, is one example, and it forms in internal combustion engines when oxygen and nitrogen react at high temperature. To draw the Lewis structure for NO, first count the valence electrons. Nitrogen contributes 5 and oxygen contributes 6, for a total of 11.

The odd electron means not every atom can reach an octet. A single N–O bond can be drawn first, and the remaining electrons are placed as lone pairs on the terminal atom. In NO, oxygen is the more electronegative atom, so it gets a full valence shell first. Because no electrons remain for a central atom, the next step is to rearrange electrons to make multiple bonds where possible. In NO, one lone pair from oxygen forms an N=O double bond, which brings nitrogen closer to an octet. A triple bond is not used, because that would give nitrogen nine electrons.

Electron-deficient molecules are another exception. In these compounds, the central atom has fewer electrons than needed for a noble-gas configuration. This often happens with central atoms from groups 2 and 13, especially when the outer atoms are hydrogen or other atoms that do not form multiple bonds. Beryllium dihydride, BeH2, and boron trifluoride, BF3, are examples. In their Lewis structures, beryllium has only four electrons and boron has only six.

BF3 can be drawn with a boron-fluorine double bond, but experimental bond lengths are closer to B–F single bonds. That evidence supports the best Lewis structure as three B–F single bonds with electron-deficient boron. The compound is also very reactive for the same reason. Boron in BF3 readily combines with a molecule that has a lone pair, such as NH3, because the lone pair on nitrogen can be shared with boron.

Hypervalent molecules are the third group. These molecules have a central atom with more than eight electrons in its valence shell. Second-period elements, with only four valence orbitals, can hold only eight electrons. Third-period and higher elements can share more than four pairs of electrons because they have empty d orbitals in the same shell. Examples include PCl5 and SF6, where phosphorus shares five pairs of electrons and sulfur shares six pairs.

Some hypervalent molecules, such as IF5 and XeF4, also have lone pairs on the central atom. In these Lewis structures, electrons remain after the outer atoms have filled octets, and those extra electrons are placed on the central atom. Together, odd-electron, electron-deficient, and hypervalent molecules show the main ways Lewis structures can depart from the octet rule.

Octet Rule Exceptions in Lewis Structures

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Learn why some Lewis structures break the octet rule with odd-electron, electron-deficient, and hypervalent molecules like NO, BF3, and SF6.

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Octet Rule Exceptions in Lewis Structures

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