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

La regla del octeto explica enlaces químicos en los compuestos del grupo principal prediciendo que cada átomo alcanza una configuración de 8 electrones. Sin embargo, existen tres excepciones importantes a esta regla. La primera excepción son las especies de electrones impares.La mayoría de las moléculas e iones tienen un número par de electrones. Sin embargo, ciertas moléculas, llamadas radicales, pueden tener uno o más electrones desapareados. Los radicales con un número impar de electrones desapareados no pueden formar octetos.El anión superóxido, un radical con un electrón desapareado, tiene 13 electrones de valencia. Puede estar representado por dos estructuras contribuyentes donde un oxígeno tiene solo siete electrones y, por lo tanto, no puede formar un octeto. La segunda excepción son los átomos que forman un octeto incompleto.Por ejemplo, el hidrógeno, el helio y el litio tienden a formar un dúo, mientras que los elementos del grupo 2 y 13 como el berilio y el boro a menudo forman moléculas con cuatro y seis electrones alrededor de ellos, respectivamente. Considere el cloruro de aluminio, que tiene 24 electrones de valencia. Mientras que todos los átomos de cloro forman un octeto, el aluminio obtiene solo 6 electrones de valencia:un octeto incompleto.Aunque el cloruro de aluminio es estable, reacciona con moléculas como el amoniaco que tienen un par de electrones no compartidos. El nitrógeno en el amoníaco dona su único par al aluminio, formando un enlace especial llamado enlace covalente coordinado, o dativo. La tercera excepción son los elementos que puede acomodar más de 8 electrones de valencia, o un octeto expandido.Estos elementos se encuentran en la tercera fila de la tabla periódica o debajo. Elementos como el fósforo, el azufre o el yodo, tienen acceso a los orbitales d, que les permite acomodar más de 8 electrones de valencia, a menudo hasta 12 o 14. Considere el anión tetracloroyoduro, que tiene 36 electrones de valencia.Incluso después de asignar los pares de electrones de enlace y satisfacer el octeto para todos los átomos, 4 electrones de valencia permanecen sin asignar. Estos electrones se colocan en el átomo de yodo central, produciendo un octeto expandido con 12 electrones. Las moléculas con más de 8 electrones de valencia alrededor del átomo central se denominan hipervalentes.Recuerde, los elementos de la segunda fila de la tabla periódica, como el carbono o el oxígeno, tienen solo orbitales s y p y nunca forman compuestos hipervalentes, porque colectivamente solo pueden contener hasta 8 electrones de valencia.

excepciones a la regla del octeto

Muchas moléculas covalentes tienen átomos centrales que no tienen ocho electrones en sus estructuras de Lewis. Estas moléculas se dividen en tres categorías:
 
 <ol>
 <li>Las moléculas de electrones impares tienen un número impar de electrones de valencia y, por lo tanto, tienen un electrón no apareado.</li>
 <li>Las moléculas con deficiencia de electrones tienen un átomo central con menos electrones de los necesarios para una configuración de gas noble.</li>
 <li>Las moléculas hipervalentes tienen un átomo central que tiene más electrones de los necesarios para una configuración de gas noble.</li>
 </ol>
 
 <h4>Moléculas de electrones impares</h4>
 
 Las moléculas que contienen un número impar de electrones se llaman radicales. El óxido nítrico, NO, es un ejemplo de una molécula de electrones impares; se produce en motores de combustión interna cuando el oxígeno y el nitrógeno reaccionan a altas temperaturas.
 
 Para dibujar la estructura de Lewis para una molécula de electrones impares como el NO, se consideran los siguientes pasos:
 
 <ol>
 <li>Determine el número total de electrones de valencia (capa externa). La suma de los electrones de valencia es 5 (de N) + 6 (de O) = 11. El número impar indica que es un radical libre, donde no todos los átomos tiene ocho electrones en su capa de valencia.</li>
 <li>Dibuje una estructura de esqueleto de la molécula. Una estructura de esqueleto con un enlace sencillo N–O puede ser dibujada fácilmente.</li>
 <li>Distribuya los electrones restantes como pares solitarios en los átomos terminales. En este caso, no hay átomo central, por lo que los electrones se distribuyen alrededor de ambos átomos. Ocho electrones son asignados al átomo más electronegativo en estas situaciones; así, el oxígeno tiene la capa de valencia llena: 
 <img alt="Figure1" src="https://www.jovecore.com/CDNSource/lessons/1066/1066Image_1.svg" style="height: 40px; display: block; margin: 0 auto;" /></li>
 <li>Coloque todos los electrones restantes en el átomo central. Dado que no quedan electrones, este paso no se aplica.</li>
 <li>Reorganice los electrones para hacer múltiples enlaces con el átomo central y obtener octetos en donde sea posible. Aunque una molécula de electrones impares no puede tener un octeto para cada átomo, cada átomo debe obtener electrones lo más cerca posible de un octeto. En este caso, el nitrógeno sólo tiene cinco electrones a su alrededor. Para acercarse a un octeto de nitrógeno, uno de los pares solitarios del oxígeno se utiliza para formar un enlace NO doble. (No puede tomarse otro par de electrones solitarios del oxígeno para formar un enlace triple porque el nitrógeno entonces tendría nueve electrones:) 
 <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>Moléculas con deficiencia de electrones</h4>
 
 Algunas moléculas, sin embargo, contienen átomos centrales que no tienen una capa de valencia llena. Generalmente, son moléculas con átomos centrales de los grupos 2 y 13, átomos externos que son hidrógeno u otros átomos que no forman enlaces múltiples. Por ejemplo, en las estructuras de Lewis del dihidruro de berilio, BeH2, y del trifluoruro de boro, BF3, los átomos de berilio y boro tienen cada uno sólo cuatro y seis electrones, respectivamente. Es posible dibujar una estructura con un enlace doble entre un átomo de boro y un átomo de flúor en BF3, satisfaciendo la regla del octeto, pero la evidencia experimental indica que las longitudes de enlace son más cercanas a las esperadas para los enlaces simples B–F. Esto sugiere que la mejor estructura de Lewis tiene tres enlaces sencillos B–F y un boro con deficiencia de electrones. La reactividad del compuesto también es consistente con el boro teniendo deficiencia de electrones. Sin embargo, los enlaces B–F son ligeramente más cortos de lo que se espera en realidad para los enlaces sencillos B–F, lo que indica que se encuentra algún carácter de enlace doble en la molécula real.
 
 <img alt="Figure3" src="https://www.jovecore.com/CDNSource/lessons/1066/1066Image_3.svg" style="height: 80px;" />
 
 Un átomo como el átomo de boro en BF3, que no tiene ocho electrones, es muy reactivo. Se combina fácilmente con una molécula que contiene un átomo con un par solitario de electrones. Por ejemplo, NH3 reacciona con BF3 porque el par solitario del nitrógeno se puede compartir con el átomo de boro:
 
 <img alt="Figure4" src="https://www.jovecore.com/CDNSource/lessons/1066/1066Image_4.svg" style="height: 80px;" />
 
 <h4>Moléculas hipervalentes</h4>
 
 Los elementos del segundo período de la tabla periódica (n = 2) pueden acomodar sólo ocho electrones en sus orbitales de la capa de valencia porque sólo tienen cuatro orbitales de valencia (uno 2s y tres orbitales 2p). Los elementos en los períodos tercero y superior (n ≥ 3) tienen más de cuatro orbitales de valencia y pueden compartir más de cuatro pares de electrones con otros átomos porque tienen orbitales d vacíos en la misma capa. Las moléculas formadas a partir de estos elementos se llaman a veces moléculas hipervalentes, como el PCl5 y el SF6. En el PCl5, el átomo central, el fósforo, comparte cinco pares de electrones. En SF6, el azufre comparte seis pares de electrones.
 
 <img alt="Figure5" src="https://www.jovecore.com/CDNSource/lessons/1066/1066Image_5.svg" style="height: 110px;" />
 
 En algunas moléculas hipervalentes, como el IF5 y el XeF4, algunos de los electrones en la capa exterior del átomo central son pares solitarios:
 
 <img alt="Figure6" src="https://www.jovecore.com/CDNSource/lessons/1066/1066image_6.svg" style="height: 110px;" />
 
 En las estructuras de Lewis para estas moléculas, quedan electrones después de llenar las capas de valencia de los átomos externos con ocho electrones. Estos electrones adicionales deben ser asignados al átomo central.
 
 Este texto es adaptado de<a href="https://openstax.org/books/chemistry-2e/pages/7-3-lewis-symbols-and-structures">Openstax, Química 2e, Sección 7.3: Símbolos y estructuras de Lewis.</a>

Excepciones a la Regla del Octeto

Muchas moléculas covalentes tienen átomos centrales que no tienen ocho electrones en sus estructuras de Lewis. Estas moléculas se dividen en tres categorías:

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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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