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

Les structures de Lewis sont des représentations simplifiées des liaisons chimiques entre les atomes. L'écriture de structures de Lewis pour des composés neutres comme l'éthène ou l'ammoniac, ou pour des ions polyatomiques tels que l'hydronium ou le phosphate implique la séquence d'étapes suivante. Tout d'abord, calculez le nombre total d'électrons de valence dans la molécule.Considérons l'éthène, par exemple. Chaque atome de carbone possède 4 électrons de valence, et chaque atome d'hydrogène possède 1 électron de valence;avec 2 atomes de carbone et 4 atomes d'hydrogène, l'éthène possède un total de 12 électrons de valence. Ensuite, dessinez la structure squelettique de la molécule.Généralement, le premier atome de la formule chimique ou l'atome moins électronégatif est placé au centre. Cet atome est aussi appelé l'atome central et doit être capable de former au moins deux obligations. Des atomes d'hydrogène ou plus d'atomes électronégatifs sont positionnés en phase terminale autour de l'atome central.Ici, les atomes de carbone occupent les positions centrales avec des atomes d'hydrogène qui les entourent. L'étape suivante est l'achèvement de l'octet. Tout d'abord, placez une paire d'électrons entre tous les deux atomes pour représenter les paires de liaison.Les électrons restants sont distribués sous forme de paires solitaires aux atomes terminaux et centraux pour satisfaire les configurations de duo ou d'octet. Dans ce cas, puisque l'hydrogène a atteint un duo, les électrons restants sont transférés aux atomes de carbone. Maintenant, utilisez les paires solitaires pour former de multiples liaisons entre les atomes qui n'ont pas d'octet.Dans l'éthène, les atomes de carbone n'ont pas encore atteint un octet. Par conséquent, la paire solitaire est déplacée vers la région de liaison pour former une double liaison permettant à tous les atomes d'atteindre une configuration électronique stable. Pour écrire la structure de Lewis pour les ions polyatomiques chargés comme l'hydronium ou le phosphate, la même approche est suivie, par une modification pour tenir compte de la charge ionique.Pour dessiner la structure de Lewis de l'hydronium, tout d'abord, calculer le nombre d'électrons de valence sur l'ion. Puisque l'ion hydronium a une charge positive, retirez un électron des électrons de valence totale, réduisant le nombre à 8. Ensuite, illustrez la molécule par sa structure squelettique.Maintenant, placez une paire électron entre chaque atome, suivie par la seule paire d'électrons sur l'atome central. De cette façon, les atomes d'hydrogéne et d'oxygène satisfont respectivement le duo et l'octet. Enfin, écrivez la structure de Lewis entre parenthèses avec une charge positive sur son coin supérieur droit.

structures de lewis des composés moléculaires et des ions polyatomiques

Afin de dessiner des structures de Lewis pour les molécules complexes et les ions moléculaires, il est utile de suivre une procédure étape par étape comme décrite : 
 
 <ol>
 <li>Déterminer le nombre total d'électrons de valence (couche externe). Pour les cations, soustraire un électron pour chaque charge positive. Pour les anions, additionnez un électron pour chaque charge négative. </li>
 <li>Dessinez la structure du squelette de la molécule ou de l'ion, en organisant les atomes autour d'un atome central (en général, l'élément le moins électronégatif doit être placé au centre). Connectez chaque atome à l'atome central avec une liaison unique (un doublet électronique). </li>
 <li>Répartissez les électrons restants sous forme de doublets libres sur les atomes terminaux (sauf l'hydrogène), en formant un octet autour de chaque atome. </li>
 <li>Placez tous les électrons restants sur l'atome central. </li>
 <li>Réorganisez les électrons des atomes externes pour créer des liaisons multiples avec l'atome central afin d'obtenir des octets dans la mesure du possible. </li>
 </ol>
 
 Par exemple, considérez SiH4, CHO2–, NO+ et OF2 comme des exemples pour lesquels cette directive générale peut être appliquée afin de déterminer leurs structures de Lewis. 
 
 <ol>
 <li>Déterminer le nombre total d'électrons de valence (couche externe) dans la molécule ou l'ion.  </li>
 </ol>
 
 Pour une molécule comme SiH4, on additionne le nombre d'électrons valence sur chaque atome de la molécule :  
 
 = [4 e−valence/atome Si × 1 atome Si] + [ 1 e−valence/atome H × 4 atomes H] = 8 e−valence
 
 <h4>CHO2–</h4>
 
 Pour un ion négatif, tel que CHO2–, le nombre d'électrons de valence sur les atomes est additionné au nombre de charges négatives sur l'ion (un électron est gagné pour chaque charge négative) : 
 
 = [4 e−valence/atome C × 1 atome C] + [1 e−valence/atome H × 1 atome H] + [6 e−valence/atome O × 2 atomes O] + [1 e−supplémentaire] = 18 e−valence
 
 <h4>NO+</h4>
 
 Pour un ion positif, tel que NO+, on additionne le nombre d'électrons de valence sur les atomes de l'ion, suivi de la soustraction du nombre de charges positives sur l'ion (un électron est perdu pour chaque charge positive) du nombre total d'électrons de valence : 
 
 = [5 e−valence/atome N × 1 atome N] + [6 e−valence/atome O × 1 atome O] + [−1 e−] = 10 e−valence
 
 <h4>OF2 </h4>
 
 OF2 étant une molécule neutre, on additionne simplement le nombre d'électrons de valence : 
 
 = [6 e−valence/atome O × 1 atome O] + [7 e&minusvalence/atome F × 2 atomes F] = 20 e−valence
 
 <ol start="2">
 <li>Tracez la structure du squelette de la molécule ou de l'ion, en arrangeant les atomes autour d'un atome central et en reliant chaque atome à l'atome central avec une liaison (un doublet électronique). (Notez que les ions sont indiqués par des crochets autour de la structure, et que la charge ionique est à l'extérieur des crochets :) 
 <img alt="Figure1" src="https://www.jovecore.com/CDNSource/lessons/1063/1063Image_1.svg" style="height: 80px;" /> 
 Dans les cas où plusieurs dispositions d'atomes sont possibles, comme pour CHO2−, on se sert des preuves expérimentales pour choisir la bonne. En général, les éléments moins électronégatifs sont plus susceptibles d'être des atomes centraux. Dans  CHO2−,  l'atome de carbone le moins électronégatif occupe la position centrale avec les atomes d'oxygène et d'hydrogène qui l'entourent. D'autres exemples incluent le P dans POCl3, le S dans SO2 et le Cl dans  ClO4−.  Une exception est que l'hydrogène n'est presque jamais un atome central. En tant qu'élément le plus électronégatif, le fluor ne peut pas non plus être un atome central. </li>
 <li>Répartissez les électrons restants comme des doublet libres sur les atomes terminaux (sauf l'hydrogène) pour compléter leurs couches de valence avec un octet d'électrons (sans électrons restants sur SiH4, sa structure ne change pas). 
 <img alt="Figure2" src="https://www.jovecore.com/CDNSource/lessons/1063/1063image_2.svg" style="height: 80px;" /></li>
 <li>Placez tous les électrons restants sur l'atome central.  
 <ul>
 <li>Pour SiH4,  CHO2−  et NO+, il n'y a pas d'électrons restants. Pour OF2, sur les 16 électrons restants, 12 sont placés, ce qui laisse ainsi 4 électrons à placer sur l'atome central : 
 <img alt="Figure3" src="https://www.jovecore.com/CDNSource/lessons/1063/1063Image_3.svg" style="height: 40px;" /></li>
 </ul>
 </li>
 <li>Réorganisez les électrons des atomes externes pour créer des liaisons multiples avec l'atome central, afin d'obtenir des octets dans la mesure du possible.  
 <ul>
 <li>SiH4 : Si a déjà un octet, donc il ne faut rien faire.  
 <img alt="Figure4" src="https://www.jovecore.com/CDNSource/lessons/1063/1063Image_4.svg" style="height: 70px;" /></li>
 <li>CHO2− :  les électrons de valence sont répartis comme des doublets libres sur les atomes d'oxygène, mais il manque un octet sur l'atome de carbone.   
 <img alt="Figure5" src="https://www.jovecore.com/CDNSource/lessons/1063/1063Image_5.svg" style="height: 80px;" /></li>
 <li>Par conséquent, un doublet libre d'électrons est donné de l'un des oxygène à l'atome de carbone, formant une liaison double. Selon l'atome d'oxygène qui a donné les électrons, il peut y avoir deux structures possibles, qui sont aussi appelées structures de résonance.  </li>
 <li>NO+ : pour cet ion, huit électrons de valence sont additionnés, mais aucun atome n'a d'octet. Des électrons supplémentaires ne peuvent pas être ajoutés car le nombre total d'électrons est déjà utilisé. Dans ce scénario, les électrons doivent être déplacés pour former des liaisons multiples. L'atome d'azote a deux doublets d'électrons libres et l'atome d'oxygène en a un.   
 <img alt="Figure6" src="https://www.jovecore.com/CDNSource/lessons/1063/1063Image_6.svg" style="height: 55px;" /></li>
 <li>Cela ne produit toujours pas d'octet, il faut donc déplacer un autre doublet pour former une liaison triple. 
  <img alt="Figure7" src="https://www.jovecore.com/CDNSource/lessons/1063/1063Image_7.svg" style="height: 50px;" /></li>
 <li>Dans OF2, rien ne change car chaque atome a déjà un octet. 
 <img alt="Figure8" src="https://www.jovecore.com/CDNSource/lessons/1063/1063Image_8.svg" style="height: 40px;" /></li>
 </ul>
 </li>
 </ol>
 
 Ce texte est adapté de <a href="https://openstax.org/books/chemistry-2e/pages/7-3-lewis-symbols-and-structures">Openstax, Chimie 2e, Section 7.3 : Symboles et structures de Lewis.</a>

Structures de Lewis des composés moléculaires et des ions polyatomiques

Pour dessiner des structures de Lewis pour des molécules complexes et des ions moléculaires, il est utile de suivre une procédure étape par étape comme décrit :

Déterminez le ...

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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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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...

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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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