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

A regra do octeto explica a ligação química nos compostos do grupo principal prevendo que cada átomo alcance uma configuração de 8-elétrons. No entanto, existem três grandes exceções a esta regra. A primeira exceção são as espécies de elétrodos estranhos.A maioria das moléculas e íons têm um número par de elétrons. No entanto, certas moléculas, chamadas radicais, ter um ou mais elétrons não emparelhados. Os radicais com um número ímpar de elétrons não pareados não podem conseguir octetos.O ânion superóxido, um radical com um elétron não emparelhado, tem 13 valências de elétrons. Pode ser representado por duas estruturas contribuintes onde um oxigênio tem apenas sete elétrons e por isso não pode alcançar um octeto. A segunda exceção são os átomos que formam um octeto incompleto.Por exemplo, hidrogênio, hélio e lítio tendem a atingir um dueto, enquanto os elementos do grupo 2 e 13 como o berílio e o boro muitas vezes formar moléculas com quatro e seis elétrons à sua volta, respectivamente. Considere o cloreto de alumínio, que tem 24 valências de elétrons. Enquanto todos os átomos de cloro alcançam o octeto, o alumínio recebe apenas 6 valências de elétrons um octeto incompleto.Embora o cloreto de alumínio seja estável, reage com moléculas como o amoníaco que têm um par de elétrons não partilhados. O nitrogênio em amoníaco doa o seu par único ao alumínio, formando uma ligação especial chamada covalente coordenada, ou dative, ligação. A terceira exceção são elementos que pode acomodar mais de 8 valências de elétrons, ou um octeto expandido.Estes elementos estão localizados na terceira fila da tabela periódica e abaixo. Elementos, tais como fósforo, enxofre, ou iodo, ter acesso a orbitais d, permitindo-lhes acomodar mais de 8 valências de elétrons muitas vezes até 12 ou 14. Considere o ânion tetracloroiodida, que tem 36 valências de elétrons.Mesmo após a atribuição dos pares de elétrons de ligação e satisfazendo o octeto para todos os átomos, 4 valências de elétrons permanecem por atribuir. Estes elétrons são colocados sobre o átomo central de iodo, produzindo um octeto expandido com 12 elétrons. Moléculas com mais de 8 valências de elétrons em torno do átomo central são chamados hipervalentes.Lembre-se, elementos da segunda fila da tabela periódica, como o carbono ou o oxigênio, ter apenas s e p orbitais e nunca formar hipervalentes compostos, porque coletivamente só podem aguentar até 8 valências de elétrons.

excepções à regra do octeto

Muitas moléculas covalentes têm átomos centrais que não têm oito eletrões nas suas estruturas de Lewis. Estas moléculas dividem-se em três categorias:

<ol>
<li>Moléculas com eletrões ímpares têm um número ímpar de eletrões de valência e, por conseguinte, têm um eletrão não emparelhado.</li>
<li>Moléculas deficientes em eletrões têm um átomo central com menos eletrões do que o necessário para uma configuração de gás nobre.</li>
<li>Moléculas hipervalentes têm um átomo central que tem mais eletrões do que o necessário para uma configuração de gás nobre.</li>
</ol>

<h4>Moléculas com Eletrões Ímpares</h4>

Moléculas que contêm um número ímpar de eletrões são chamadas radicais. O óxido nítrico, NO, é um exemplo de uma molécula com eletrões ímpares; é produzido em motores de combustão interna quando o oxigénio e o azoto reagem a temperaturas elevadas.

Para desenhar a estrutura de Lewis para uma molécula com eletrões ímpares como o NO, são considerados os seguintes passos:

<ol>
<li>Determine o número total de eletrões de valência (camada exterior). A soma dos eletrões de valência é 5 (do N) + 6 (do O) = 11. O número ímpar indica que é um radical livre, onde nem todos os átomos têm oito eletrões na sua camada de valência.</li>
<li>Desenhe uma estrutura de esqueleto da molécula. Uma estrutura de esqueleto com uma ligação simples N–O pode ser facilmente desenhada.</li>
<li>Distribua os eletrões restantes como pares solitários nos átomos terminais. Neste caso, não existe um átomo central, portanto os eletrões são distribuídos em torno de ambos os átomos. Oito eletrões são atribuídos ao átomo mais eletronegativo nestas situações; assim, o oxigénio tem a camada de valência preenchida: 
<img alt="Figure1" src="https://www.jove.com/files/core/CDNSource/lessons/1066/1066Image_1.svg" style="height: 40px; display: block; margin: 0 auto;" /></li>
<li>Coloque todos os eletrões restantes no átomo central. Uma vez que não existem eletrões restantes, este passo não se aplica.</li>
<li>Reorganize os eletrões para fazer múltiplas ligações com o átomo central, de forma a obter octetos sempre que possível. Embora uma molécula com eletrões ímpares não possa ter um octeto para cada átomo, cada átomo deve obter eletrões o mais próximo possível de um octeto. Neste caso, o azoto tem apenas cinco eletrões à sua volta. Para se aproximar de um octeto para o azoto, um dos pares solitários de oxigénio é utilizado para formar uma ligação NO dupla. (Outro par isolado de eletrões não pode ser retirado do oxigénio para formar uma ligação tripla porque o azoto ficaria com nove eletrões:) 
<img alt="Figure2" src="https://www.jove.com/files/core/CDNSource/lessons/1066/1066Image_2.svg" style="height: 40px; display: block; margin: 0 auto;" /></li>
</ol>

<h4>Moléculas Deficientes em Eletrões</h4>

Algumas moléculas, no entanto, contêm átomos centrais que não têm uma camada de valência preenchida. Geralmente, são moléculas com átomos centrais dos grupos 2 e 13, átomos exteriores de hidrogénio, ou outros átomos que não formam ligações múltiplas. Por exemplo, nas estruturas de Lewis do dihidreto de berílio, BeH2 e trifluoreto de boro, BF3, os átomos de berílio e boro têm apenas quatro e seis eletrões, respectivamente. É possível desenhar uma estrutura com dupla ligação entre um átomo de boro e um átomo de flúor em BF3, satisfazendo a regra de octeto, mas evidências experimentais indicam que os comprimentos de ligação estão mais próximos do esperado para ligações simples B–F. Isto sugere que a melhor estrutura de Lewis tem três ligações simples B–F e boro deficiente em eletrões. A reatividade do composto também é consistente com boro deficiente em eletrões. No entanto, ligações B–F são ligeiramente mais curtas do que o que é realmente esperado para ligações simples B–F, indicando que algum carácter de ligação dupla é encontrado na molécula real.

<img alt="Figure3" src="https://www.jove.com/files/core/CDNSource/lessons/1066/1066Image_3.svg" style="height: 80px;" />

Um átomo como o átomo de boro em BF3, que não tem oito eletrões, é muito reativo. Combina-se prontamente com uma molécula que contém um átomo com um par solitário de eletrões. Por exemplo, NH3 reage com BF3 porque o par isolado no azoto pode ser partilhado com o átomo de boro:

<img alt="Figure4" src="https://www.jove.com/files/core/CDNSource/lessons/1066/1066Image_4.svg" style="height: 80px;" />

<h4>Moléculas Hipervalentes</h4>

Elementos no segundo período da tabela periódica (n = 2) podem apenas acomodar oito eletrões nas respectivas orbitais da camada de valência porque têm apenas quatro orbitais de valência (uma orbital 2s e três 2p). Os elementos no terceiro período e superiores (n ≥ 3) têm mais de quatro orbitais de valência e podem compartilhar mais de quatro pares de eletrões com outros átomos porque têm orbitais d vazias na mesma camada. As moléculas formadas a partir desses elementos são às vezes chamadas moléculas hipervalentes, como PCl5, e SF6. Em PCl5, o átomo central, fósforo, compartilha cinco pares de eletrões. Em SF6, o enxofre compartilha seis pares de eletrões.

<img alt="Figure5" src="https://www.jove.com/files/core/CDNSource/lessons/1066/1066Image_5.svg" style="height: 110px;" />

Em algumas moléculas hipervalentes, como IF5 e XeF4, alguns dos eletrões na camada exterior do átomo central são pares solitários:

<img alt="Figure6" src="https://www.jove.com/files/core/CDNSource/lessons/1066/1066image_6.svg" style="height: 110px;" />

Nas estruturas de Lewis para estas moléculas, sobram eletrões depois de preencher as camadas de valência dos átomos exteriores com oito eletrões. Estes eletrões adicionais devem ser atribuídos ao átomo central.

Este texto é adaptado de <a href="https://openstax.org/books/chemistry-2e/pages/7-3-lewis-symbols-and-structures">Openstax, Chemistry 2e, Section 7.3: Lewis Symbols and Structures.</a>

Excepções à Regra do Octeto

Muitas moléculas covalentes têm átomos centrais que não possuem oito elétrons em suas estruturas de Lewis. Essas moléculas se enquadram em três categorias:

...

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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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Octet Rule Exceptions in Lewis Structures

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