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

La regola dell'ottetto spiega il legame chimico nei composti del gruppo principale, prevedendo che ogni atomo raggiunga una configurazione di 8 elettroni. Tuttavia, ci sono tre principali eccezioni a questa regola. La prima eccezione è rappresentata dalle specie di elettroni dispari.

La maggior parte delle molecole e degli ioni ha un numero pari di elettroni. Tuttavia, alcune molecole, chiamate radicali, hanno uno o più elettroni spaiati. I radicali con un numero dispari di elettroni spaiati non possono ottenere ottetti.

L'anione superossido, un radicale con un elettrone spaiato, ha 13 elettroni di valenza. Può essere rappresentato da due strutture contribuenti, in cui un ossigeno ha solo sette elettroni e quindi non può raggiungere un ottetto. La seconda eccezione sono gli atomi che formano un ottetto incompleto.

Per esempio, idrogeno, elio e litio tendono a raggiungere un duetto, mentre elementi del gruppo 2 e 13, come il berillio e il boro, spesso formano molecole con quattro e sei elettroni intorno, rispettivamente. Considerate il cloruro di alluminio, che ha 24 elettroni di valenza. Mentre tutti gli atomi di cloro raggiungono l'ottetto, l'alluminio riceve solo 6 elettroni di valenza, ossia un ottetto incompleto.

Benché il cloruro di alluminio sia stabile, reagisce con molecole, come l'ammoniaca, che hanno una coppia di elettroni non condivisa. L'azoto nell'ammoniaca dona la sua coppia solitaria all'alluminio, formando un legame speciale detto legame covalente coordinato, o dativo. La terza eccezione sono gli elementi che possono ospitare più di 8 elettroni di valenza, o un ottetto espanso.

Questi elementi si trovano nella terza riga della tavola periodica e più in basso. Elementi, come fosforo, zolfo o iodio, hanno accesso agli orbitali d, consentendo loro di ospitare più di 8 elettroni di valenza, spesso fino a 12 o 14. Considerate l'anione tetracloroioduro, che ha 36 elettroni di valenza.

Anche dopo aver assegnato le coppie di elettroni di legame e soddisfatto l'ottetto per tutti gli atomi, 4 elettroni di valenza rimangono non assegnati. Questi elettroni sono posti sull'atomo di iodio centrale, producendo un ottetto espanso con 12 elettroni. Le molecole con più di 8 elettroni di valenza attorno all'atomo centrale sono chiamate ipervalenti.

Ricordate, gli elementi della seconda riga della tavola periodica, come il carbonio o l'ossigeno, hanno solo orbitali s e p e non formano mai composti ipervalenti, perché collettivamente possono contenere solo fino a 8 elettroni di valenza.

eccezioni alla regola dell'ottetto

Molte molecole covalenti hanno atomi centrali che non hanno otto elettroni nelle loro strutture di Lewis. Queste molecole rientrano in tre categorie:

<ol>
	<li>Le molecole con elettroni dispari hanno un numero dispari di elettroni di valenza e quindi hanno un elettrone spaiato.</li>
	<li>Le molecole carenti di elettroni hanno un atomo centrale con meno elettroni di quelli necessari per una configurazione di gas nobile.</li>
	<li>Le molecole ipervalenti hanno un atomo centrale che ha più elettroni del necessario per una configurazione di gas nobile.</li>
</ol>

Molecole di elettroni dispari

Le molecole che contengono un numero dispari di elettroni sono chiamate radicali. L'ossido nitrico, NO, è un esempio di molecola di elettroni dispari; viene prodotto nei motori a combustione interna quando ossigeno e azoto reagiscono ad alte temperature.

Per disegnare la struttura di Lewis per una molecola di elettroni dispari come NO, vengono considerati i seguenti passaggi:

<ol>
	<li>Determinare il numero totale di elettroni di valenza (guscio esterno). La somma degli elettroni di valenza è 5 (da N) + 6 (da O) = 11. Il numero dispari indica che si tratta di un radicale libero, dove non tutti gli atomi hanno otto elettroni nel suo guscio di valenza.</li>
	<li>Disegna una struttura scheletrica della molecola. È possibile disegnare facilmente una struttura a scheletro con un singolo legame N–O.</li>
	<li>Distribuire gli elettroni rimanenti come coppie solitarie sugli atomi terminali. In questo caso non esiste un atomo centrale, quindi gli elettroni sono distribuiti attorno a entrambi gli atomi. In queste situazioni otto elettroni vengono assegnati all'atomo più elettronegativo; quindi, l'ossigeno ha il guscio di valenza pieno:
	<img alt="Figure1" src="https://www.jove.com/files/core/CDNSource/lessons/1066/1066Image_1.svg" style="height: 40px;" />
	</li>
	<li>Posiziona tutti gli elettroni rimanenti sull'atomo centrale. Visto che non ci sono elettroni rimanenti, questo passaggio non è applicabile.</li>
	<li>Riorganizzare gli elettroni per creare legami multipli con l'atomo centrale al fine di ottenere ottetti ove possibile. Sebbene una molecola con elettroni dispari non possa avere un ottetto per ogni atomo, ogni atomo dovrebbe avere elettroni il più vicino possibile a un ottetto. In questo caso, l'azoto ha solo cinque elettroni attorno a sé. Per avvicinarsi a un ottetto per l'azoto, una delle coppie solitarie dell'ossigeno viene utilizzata per formare un doppio legame NO. (Un’altra coppia solitaria di elettroni non può essere presa dall’ossigeno per formare un triplo legame perché l’azoto avrebbe allora nove elettroni :)
	<img alt="Figure2" src="https://www.jove.com/files/core/CDNSource/lessons/1066/1066Image_2.svg" style="height: 40px;" />
	</li>
</ol>

Molecole carenti di elettroni

Alcune molecole, tuttavia, contengono atomi centrali che non hanno un guscio di valenza pieno. Generalmente si tratta di molecole con atomi centrali dei gruppi 2 e 13, atomi esterni che sono idrogeno o altri atomi che non formano legami multipli. Ad esempio, nelle strutture di Lewis del diidruro di berillio, BeH_2, e del trifluoruro di boro, BF_3, gli atomi di berillio e boro hanno ciascuno solo quattro e sei elettroni, rispettivamente. È possibile disegnare una struttura con un doppio legame tra un atomo di boro e un atomo di fluoro in BF_3, soddisfacendo la regola dell'ottetto ma, prove sperimentali indicano che le lunghezze dei legami sono più vicine a quelle previste per i singoli legami B-F. Ciò suggerisce che la migliore struttura di Lewis ha tre legami singoli B – F e boro carente di elettroni. La reattività del composto è coerente anche con il boro carente di elettroni. Tuttavia, i legami B–F sono leggermente più corti di quanto effettivamente previsto per i singoli legami B–F, indicando che nella molecola reale si trova qualche carattere di doppio legame.

<img alt="Figure3" src="https://www.jove.com/files/core/CDNSource/lessons/1066/1066Image_3.svg" style="height: 80px;" />

Un atomo come l'atomo di boro in BF3, che non ha otto elettroni, è molto reattivo. Si combina facilmente con una molecola contenente un atomo con una coppia solitaria di elettroni. Ad esempio, NH_3 reagisce con BF_3 perché la coppia solitaria sull'azoto può essere condivisa con l'atomo di boro:

<img alt="Figure4" src="https://www.jove.com/files/core/CDNSource/lessons/1066/1066Image_4.svg" style="height: 80px;" />

Molecole ipervalenti

Gli elementi nel secondo periodo della tavola periodica (n = 2) possono ospitare solo otto elettroni nei loro orbitali del guscio di valenza perché hanno solo quattro orbitali di valenza (uno orbitale 2s e tre orbitali 2p). Gli elementi nel terzo periodo e in quelli superiori (n ≥ 3) hanno più di quattro orbitali di valenza e possono condividere più di quattro coppie di elettroni con altri atomi perché hanno orbitali d vuoti nello stesso livello. Le molecole formate da questi elementi sono talvolta chiamate molecole ipervalenti, come PCl_5 e SF_6. Nel PCl_5, l'atomo centrale, il fosforo, condivide cinque coppie di elettroni. In SF_6, lo zolfo condivide sei coppie di elettroni.

<img alt="Figure5" src="https://www.jove.com/files/core/CDNSource/lessons/1066/1066Image_5.svg" style="height: 110px;" />

In alcune molecole ipervalenti, come IF_5 e XeF_4, alcuni degli elettroni nel guscio esterno dell'atomo centrale sono coppie solitarie:

<img alt="Figure6" src="https://www.jove.com/files/core/CDNSource/lessons/1066/1066image_6.svg" style="height: 110px;" />

Nelle strutture di Lewis di queste molecole, ci sono elettroni rimasti dopo aver riempito i gusci di valenza degli atomi esterni con otto elettroni. Questi elettroni aggiuntivi devono essere assegnati all'atomo centrale.

Questo testo è adattato da <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>

Eccezioni alla regola dell'ottetto

Molte molecole covalenti hanno atomi centrali che non hanno otto elettroni nelle loro strutture di Lewis. Queste molecole si dividono in tre categorie:

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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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