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

Die Oktett-Regel erklärt die chemische Bindung für Hauptgruppenverbindungen durch die Vorhersage, dass jedes Atom eine 8-Elektronen-Konfiguration erreicht. Es gibt jedoch drei große Ausnahmen von dieser Regel. Die erste Ausnahme sind ungerade Elektronenarten.Die meisten Moleküle und Ionen haben eine gerade Anzahl von Elektronen. Bestimmte Moleküle, die als Radikale bezeichnet werden, haben jedoch ein oder mehrere ungepaarte Elektronen. Radikale mit einer ungeraden Anzahl ungepaarter Elektronen können keine Oktette erreichen.Das Superoxid-Anion, ein Radikal mit einem freiem Elektronpaar, hat 13 Valenzelektronen. Es kann durch zwei beitragende Strukturen dargestellt werden, wobei eine nur sieben Elektronen im Sauerstoff hat und somit kein Oktett erreichen kann. Die zweite Ausnahme sind Atome, die ein unvollständiges Oktett bilden.Zum Beispiel neigen Wasserstoff, Helium und Lithium dazu, ein Duett zu erreichen, während Elemente der 2ten und 13ten Gruppe, wie Beryllium und Bor, oft Moleküle mit vier bzw. sechs Elektronen um sich herum bilden. Betrachten Sie Aluminiumchlorid, welches 24 Valenzelektronen besitzt.Während alle Chloratome das Oktett erreichen, erhält Aluminium nur 6 Valenzelektronen ein unvollständiges Oktett. Obwohl Aluminiumchlorid stabil ist, reagiert es mit Molekülen wie Ammoniak, die ein freies Elektronenpaar haben. Der Stickstoff im Ammoniak gibt sein freies Elektronenpaar an Aluminium ab und bildet eine spezielle Bindung, die als koordinierte kovalente oder dative Bindung bezeichnet wird.Die dritte Ausnahme bilden Elemente, die mehr als 8 Valenzelektronen oder ein expandiertes Oktett aufnehmen können. Diese Elemente befinden sich in der dritten Reihe des Periodensystems und darunter. Elemente wie Phosphor, Schwefel oder Jod haben Zugang zu d-Orbitalen, sodass sie mehr als 8 Valenzelektronen, oft bis zu 12 oder 14, aufnehmen können.Betrachten Sie das Tetrachloriodid-Anion, das 36 Valenzelektronen besitzt. Selbst nach Zuweisung der Bindungselektronenpaare und Erfüllung des Oktettes für alle Atome bleiben 4 Valenzelektronen frei. Diese Elektronen werden auf das zentrale Jodatom gesetzt, was somit ein erweitertes Oktett mit 12 Elektronen ergibt.Moleküle mit mehr als 8 Valenzelektronen um das Zentralatom herum werden als hypervalent bezeichnet. Denken Sie daran, dass Elemente aus der zweiten Reihe des Periodensystems, wie z. B.Kohlenstoff oder Sauerstoff, nur s-und p-Orbitale haben und niemals hypervalente Verbindungen bilden, da sie zusammen nur bis zu 8 Valenzelektronen aufnehmen können.

ausnahmen von der oktettregel

Viele kovalente Moleküle haben Zentralatome, deren Lewis-Strukturen nicht über acht Elektronen verfügen. Diese Moleküle fallen in drei Kategorien: <ol> <li>Moleküle mit ungeraden Elektronen haben eine ungerade Anzahl an Valenzelektronen und daher ein ungepaartes Elektron.</li> <li>Elektronenarme Moleküle haben ein Zentralatom mit weniger Elektronen, als für eine Edelgaskonfiguration erforderlich sind.</li> <li>Hypervalente Moleküle haben ein Zentralatom, das mehr Elektronen hat, als für eine Edelgaskonfiguration erforderlich sind.</li> </ol> <h4>Moleküle mit ungeraden Elektronen</h4> Moleküle, die eine ungerade Anzahl an Elektronen enthalten, werden Radikale genannt. Stickstoffmonoxid, NO, ist ein Beispiel für ein Molekül mit ungeraden Elektronen; Es entsteht in Verbrennungsmotoren, wenn Sauerstoff und Stickstoff bei hohen Temperaturen reagieren. Um die Lewis-Struktur für ein Molekül mit ungeraden Elektronen wie NO zu zeichnen, sollten die folgenden Schritte berücksichtigt werden: <ol> <li>Bestimmen Sie die Gesamtzahl der Valenzelektronen (äußere Schale). Die Summe der Valenzelektronen beträgt 5 (aus N) + 6 (aus O) = 11. Die ungerade Zahl zeigt an, dass es sich um ein freies Radikal handelt, bei dem nicht jedes Atom acht Elektronen in seiner Valenzschale hat.</li> <li>Zeichnen Sie eine Skelettstruktur des Moleküls. Eine Gerüststruktur mit einer N-O-Einfachbindung lässt sich leicht zeichnen.</li> <li>Verteilen Sie die verbleibenden Elektronen als freie Elektronenpaare auf die terminalen Atome. In diesem Fall gibt es kein Zentralatom, die Elektronen sind also um beide Atome verteilt. In diesen Situationen werden dem elektronegativeren Atom acht Elektronen zugeordnet; Sauerstoff hat also die gefüllte Valenzschale: <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>Platzieren Sie alle verbleibenden Elektronen auf das Zentralatom. Da keine verbleibenden Elektronen vorhanden sind, entfällt dieser Schritt.</li> <li>Ordnen Sie die Elektronen neu an, um Mehrfachbindungen mit dem Zentralatom einzugehen und so, wo immer möglich, Oktette zu erhalten. Obwohl ein Molekül mit ungeraden Elektronen nicht für jedes Atom ein Oktett haben kann, sollte jedes Atom Elektronen so nah wie möglich an einem Oktett haben. In diesem Fall hat Stickstoff nur fünf Elektronen um sich herum. Um näher an ein Oktett für Stickstoff zu gelangen, wird eines der freien Elektronenpaare von Sauerstoff genutzt, um eine NO-Doppelbindung zu bilden. (Ein weiteres freies Elektronenpaar kann dem Sauerstoff nicht entnommen werden, um eine Dreifachbindung zu bilden, da Stickstoff dann neun Elektronen hätte:) <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>Elektronenarme Moleküle</h4> Einige Moleküle enthalten jedoch Zentralatome, die keine gefüllte Valenzschale besitzen. Im Allgemeinen handelt es sich dabei entweder um Moleküle mit zentralen Atomen der Gruppen 2 und 13, äußere Wasserstoffatome oder anderen Atome, die keine Mehrfachbindungen eingehen. Beispielsweise haben in den Lewis-Strukturen von Berylliumdihydrid (BeH2) und Bortrifluorid (BF3) die Beryllium- und Boratome jeweils nur vier bzw. sechs Elektronen. Es ist möglich, eine Struktur mit einer Doppelbindung zwischen einem Boratom und einem Fluoratom in BF3 zu zeichnen, welche die Oktettregel erfüllt, aber experimentelle Erkenntnisse deuten darauf hin, dass die Bindungslängen näher an den für die B-F-Einfachbindungen erwarteten liegen. Dies legt nahe, dass die beste Lewis-Struktur drei B-F-Einfachbindungen und ein elektronenarmes Bor als Zentrum aufweist. Die Reaktivität der Verbindung stimmt auch mit der Annahme eines elektronenarmen Boratoms überein. Allerdings sind die B-F-Bindungen etwas kürzer als eigentlich für B-F-Einfachbindungen zu erwarten wäre, was darauf hindeutet, dass im tatsächlichen Molekül doch ein gewisser Doppelbindungscharakter vorliegt. <img alt="Figure3" src="https://www.jove.com/files/core/CDNSource/lessons/1066/1066Image_3.svg" style="height: 80px;"> Ein Atom wie das Boratom in BF3, das nicht über acht Elektronen verfügt, ist sehr reaktiv. Es verbindet sich leicht mit einem Molekül, das ein Atom mit einem freien Elektronenpaar enthält. Beispielsweise reagiert NH3 mit BF3, da das freie Elektronenpaar am Stickstoff mit dem Boratom geteilt werden kann: <img alt="Figure4" src="https://www.jove.com/files/core/CDNSource/lessons/1066/1066Image_4.svg" style="height: 80px;"> <h4>Hypervalente Moleküle</h4> Elemente in der zweiten Periode des Periodensystems (n = 2) können nur acht Elektronen in ihren Valenzschalenorbitalen aufnehmen, da sie nur vier Valenzorbitale haben (ein 2s- und drei 2p-Orbitale). Elemente in der dritten und höheren Perioden (n ≥ 3) haben mehr als vier Valenzorbitale und können mehr als vier Elektronenpaare mit anderen Atomen teilen, da sie auch leere d-Orbitale in derselben Schale haben. Aus diesen Elementen gebildete Moleküle werden manchmal als hypervalente Moleküle wie PCl5 und SF6 bezeichnet. In PCl5 teilt sich das Zentralatom Phosphor fünf Elektronenpaare. In SF6 teilt sich Schwefel sechs Elektronenpaare. <img alt="Figure5" src="https://www.jove.com/files/core/CDNSource/lessons/1066/1066Image_5.svg" style="height: 110px;"> In einigen hypervalenten Molekülen wie IF5 und XeF4 sind einige der Elektronen in der Außenhülle des Zentralatoms freie Elektronenpaare: <img alt="Figure6" src="https://www.jove.com/files/core/CDNSource/lessons/1066/1066image_6.svg" style="height: 110px;"> In den Lewis-Strukturen dieser Moleküle bleiben Elektronen übrig, nachdem die Valenzschalen der äußeren Atome mit acht Elektronen gefüllt wurden. Diese zusätzlichen Elektronen müssen dem Zentralatom zugeordnet werden. Dieser Text wurde adaptiert von <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>

Ausnahmen von der Oktettregel

Viele kovalente Moleküle haben Zentralatome, die keine acht Elektronen in ihren Lewis-Strukturen haben. Diese Moleküle lassen sich in drei Kategorien einteilen:

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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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Many covalent molecules have central atoms that do not have eight electrons in their Lewis structures. These molecules fall into three categories:


	...

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PS.Chm.3.2

Analyze the structure of atoms and isotopes.

Octet Rule Exceptions in Lewis Structures

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