Name: drawing lewis structures with valence electron counts
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Lewis Structures of Molecular Compounds and Polyatomic Ions

Le strutture di Lewis sono rappresentazioni semplificate dei legami chimici tra gli atomi. Scrivere le strutture di Lewis per composti neutri, come etene o ammoniaca, o per ioni poliatomici, come idronio o fosfato, implica la seguente sequenza di passaggi:Innanzitutto, calcoliamo il numero totale di elettroni di valenza nella molecola. Considerate l'etene, per esempio.Ogni atomo di carbonio ha 4 elettroni di valenza, e ogni atomo di idrogeno ha 1 elettrone di valenza;con 2 atomi di carbonio e 4 di idrogeno, l'etene ha un totale di 12 elettroni di valenza. Quindi, si disegni la struttura scheletrica della molecola. Generalmente, il primo atomo nella formula chimica o l'atomo meno elettronegativo, è posizionato centralmente.Questo atomo è anche detto atomo centrale e deve essere in grado di formare almeno due legami. Atomi di idrogeno o atomi più elettronegativi, sono posizionati all'estremità attorno all'atomo centrale. Qui, gli atomi di carbonio occupano le posizioni centrali con atomi di idrogeno che li circondano.Il passaggio successivo è il completamento dell'ottetto. Innanzitutto, posizioniamo una coppia di elettroni ogni due atomi, per rappresentare le coppie di legame. Gli elettroni rimanenti sono distribuiti come coppie solitarie agli atomi terminali e centrali, per soddisfare le configurazioni del duetto o dell'ottetto.In questo caso, poiché l'idrogeno ha raggiunto un duetto, gli elettroni rimanenti vengono trasferiti agli atomi di carbonio. Ora, utilizziamo le coppie solitarie per formare più legami fra gli atomi che mancano di un ottetto. Nell'etene, gli atomi di carbonio non hanno ancora raggiunto un ottetto.Pertanto, la coppia solitaria viene spostata nella regione del legame per formare un doppio legame che consenta a tutti gli atomi di raggiungere una configurazione elettronica stabile. Per scrivere la struttura di Lewis per gli ioni poliatomici carichi come l'idronio o il fosfato, viene seguito lo stesso approccio, con una modifica per sistemare la carica ionica. Per disegnare la struttura di Lewis dell'idronio, prima si calcoli il numero di elettroni di valenza sullo ione.Poiché lo ione idronio ha una carica positiva, si rimuova un elettrone dagli elettroni di valenza totali, riducendo il conteggio a 8. Successivamente, si illustri la molecola dalla sua struttura scheletrica. Ora si posizioni una coppia di elettroni fra ogni atomo, seguita dalla coppia di elettroni solitaria sull'atomo centrale.In questo modo, gli atomi di ossigeno idrogenato soddisfano rispettivamente il duetto e l'ottetto. Infine, si scriva la struttura di Lewis fra parentesi con una carica positiva nell'angolo in alto a destra.

strutture di lewis di composti molecolari e ioni poliatomici

Per disegnare strutture di Lewis per molecole complicate e ioni molecolari, è utile seguire una procedura passo-passo come delineato:

<ol><li>Determinare il numero totale di elettroni di valenza (guscio esterno). Per i ations, sottrarre un elettrone per ogni carica positiva. Per gli anioni, aggiungere un elettrone per ogni carica negativa.</li>
	<li>Disegna una struttura scheletrica della molecola o dello ione, disponendo gli atomi attorno a un atomo centrale. (Generalmente, l'elemento meno elettronegativo deve essere posizionato al centro.) Collegare ogni atomo all'atomo centrale con un singolo legame (una coppia di elettroni).</li>
	<li>Distribuire gli elettroni rimanenti come coppie solitarie sugli atomi terminali (eccetto l'idrogeno), completando un ottetto attorno ad ogni atomo.</li>
	<li>Posizionare tutti gli elettroni rimanenti sull'atomo centrale.</li>
	<li>Riorganizzare gli elettroni degli atomi esterni per creare legami multipli con l'atomo centrale al fine di ottenere ottetti ove possibile.</li>
</ol>Si consideri, ad esempio, SiH4, CHO2–, NO+e OF2 come esempi per i quali questa linea guida generale può essere applicata per determinare le proprie strutture di Lewis.

<ol><li>Determinare il numero totale di elettroni di valenza (guscio esterno) nella molecola o nello ione.</li>
</ol>Per una molecola come SiH4, viene aggiunto il numero di elettroni di valenza su ogni atomo nella molecola:

= [4 valenza e−/Si atomo × 1 Atomo di Si] + [ 1 valenza e−/H atomo × 4 Atomi H] = 8 valenza e−

<h4>CHO2</h4>

Per uno ione negativo, come CHO2–, il numero di elettroni di valenza sugli atomi viene aggiunto al numero di cariche negative sullo ione (un elettrone viene guadagnato per ogni singola carica negativa):

= [4 valenza e−/C atomo × 1 atomo C] + [1 valenza e−/H atomo × 1 atomo H] + [6 valenza e−/O atomo × 2 Atomi O] + [1 eaggiuntivo −] = 18 valenza e−

<h4>NO+</h4>

Per uno ione positivo, come NO+, vengono aggiunti i numeri di elettroni di valenza sugli atomi nello ione, seguiti dalla sottrazione del numero di cariche positive sullo ione (un elettrone viene perso per ogni singola carica positiva) dal numero totale di elettroni di valenza:

= [5 valenza e−/N atomo × 1 N atomo] + [6 valenza e−/O atomo × 1 O atomo] + [−1 e−] = 10 valenza e−

<h4>DI2 </h4>

OF2 essendo una molecola neutra, il numero di elettroni di valenza viene semplicemente aggiunto:

= [6 valenza e−/O atomo × 1 Atomo O] + [7 valenza e−/F atomo × 2 Atomi F] = 20 valenza e−

<ol start="2"><li>Disegna una struttura scheletrica della molecola o dello ione, disponendo gli atomi attorno a un atomo centrale e collegando ogni atomo all'atomo centrale con un singolo legame (una coppia di elettroni). (Si noti che gli ioni sono denotati con parentesi attorno alla struttura e la carica ionica al di fuori delle parentesi:) <img alt="Figure1" src="https://www.jove.com/files/core/CDNSource/lessons/1063/1063Image_1.svg" style="height: 80px;"> Nei casi in cui sono possibili diverse disposizioni di atomi, come per CHO2−, l'evidenza sperimentale viene utilizzata per scegliere quella corretta. In generale, gli elementi meno elettronegativi hanno maggiori probabilità di essere atomi centrali. In CHO2−, l'atomo di carbonio meno elettronegativo occupa la posizione centrale con gli atomi di ossigeno e idrogeno che lo circondano. Altri esempi includono P in POCl3, S in SO2e Cl in ClO4−.  Un'eccezione è che l'idrogeno non è quasi mai un atomo centrale. Come elemento più elettronegativo, anche il fluoro non può essere un atomo centrale.</li>
	<li>Distribuire gli elettroni rimanenti come coppie solitarie sugli atomi terminali (eccetto l'idrogeno) per completare i loro gusci di valenza con un ottetto di elettroni. (Senza elettroni rimanenti su SiH4, la sua struttura è invariata.) <img alt="Figure2" src="https://www.jove.com/files/core/CDNSource/lessons/1063/1063image_2.svg" style="height: 80px;">
</li>
	<li>Posizionare tutti gli elettroni rimanenti sull'atomo centrale.<ul>
<li>Per SiH4, CHO2−e NO+, non ci sono elettroni rimanenti. Per OF2, dei 16 elettroni rimanenti, 12 sono posizionati, lasciando così 4 elettroni da posizionare sull'atomo centrale: <img alt="Figure3" src="https://www.jove.com/files/core/CDNSource/lessons/1063/1063Image_3.svg" style="height: 40px;">
</li>
	</ul>
</li>
	<li>Riorganizzare gli elettroni degli atomi esterni per creare legami multipli con l'atomo centrale, per ottenere ottetti ove possibile.<ul>
<li>SiH4: Si ha già un ottetto, quindi non c'è bisogno di fare nulla. <img alt="Figure4" src="https://www.jove.com/files/core/CDNSource/lessons/1063/1063Image_4.svg" style="height: 70px;">
</li>
		<li>CHO2−: Gli elettroni di valenza sono distribuiti come coppie solitarie sugli atomi di ossigeno, ma l'atomo di carbonio manca di un ottetto. <img alt="Figure5" src="https://www.jove.com/files/core/CDNSource/lessons/1063/1063Image_5.svg" style="height: 80px;">
</li>
		<li>Quindi, una coppia solitaria di elettroni viene donata da uno dell'ossigeno all'atomo di carbonio formando un doppio legame. A seconda di quale atomo di ossigeno ha donato gli elettroni, ci possono essere due possibili strutture, altrimenti chiamate strutture di risonanza.</li>
		<li>NO+: Per questo ione vengono aggiunti otto elettroni di valenza, ma nessuno dei due atomi ha un ottetto. Non è possibile aggiungere elettroni aggiuntivi poiché gli elettroni totali sono già utilizzati. In questo scenario, gli elettroni devono essere spostati per formare legami multipli. L'atomo di azoto ha due coppie solitarie di elettroni e l'atomo di ossigeno ne ha uno. <img alt="Figure6" src="https://www.jove.com/files/core/CDNSource/lessons/1063/1063Image_6.svg" style="height: 55px;">
</li>
		<li>Questo ancora non produce un ottetto, quindi un'altra coppia deve essere spostata per formare un triplo legame. 
		 <img alt="Figure7" src="https://www.jove.com/files/core/CDNSource/lessons/1063/1063Image_7.svg" style="height: 50px;">
</li>
		<li>In OF2, nulla cambia in quanto ogni atomo ha già un ottetto. <img alt="Figure8" src="https://www.jove.com/files/core/CDNSource/lessons/1063/1063Image_8.svg" style="height: 40px;">
</li>
	</ul>
</li>
</ol>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>

Strutture di Lewis di composti molecolari e ioni poliatomici

Per disegnare strutture di Lewis per molecole complicate e ioni molecolari, è utile seguire una procedura passo-passo come descritto:

Determina il...

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drawing lewis structures with valence electron counts

Drawing Lewis structures starts with counting valence electrons and building the atom arrangement step by step. Valence electrons are the outer-shell electrons that help show how atoms bond. For ions, a positive charge means one electron is removed for each plus sign, and a negative charge means one electron is added for each minus sign.

The next step is to sketch the skeleton structure. The least electronegative atom usually goes in the center, while hydrogen is almost never central. Fluorine also cannot be a central atom because it is the most electronegative element. Ions are written in brackets, with the charge shown outside the brackets.

After the skeleton is drawn, connect the atoms to the central atom with single bonds. Each single bond represents one electron pair. Then place the remaining electrons as lone pairs on the terminal atoms, except hydrogen, until those atoms have octets when possible. Any electrons still left go on the central atom.

If the central atom still does not have an octet, move lone pairs to make double or triple bonds. This can create resonance structures, which are different valid arrangements of electrons. For the formate ion, CHO2−, one oxygen can donate a lone pair to make a double bond with carbon. That gives two possible resonance forms.

The examples in the model show how the procedure works in real cases. SiH4 has 8 valence electrons total, so its structure does not need more changes once the single bonds are drawn. CHO2− has 18 valence electrons. NO+ has 10 valence electrons, and its atoms must be rearranged into a triple bond to satisfy octets as much as possible. OF2 has 20 valence electrons, and after the lone pairs are placed, its structure stays unchanged because each atom already has an octet.

Drawing Lewis Structures with Valence Electron Counts

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Learn the step-by-step method for drawing Lewis structures, counting valence electrons, placing lone pairs, and handling ions, resonance, and multiple bonds.

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Drawing Lewis Structures with Valence Electron Counts

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