Name: shared electrons in covalent molecules
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Chemistry

Matter and its Interactions

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Covalent Bonding and Lewis Structures

Les liaisons covalentes sont formées entre les non-métaux par le partage des électrons de valence. Mais pourquoi les non-métaux préfèrent-ils partager des électrons plutôt que de les transférer comme dans les liaisons ioniques? Les non-métaux ont des énergies d'ionisation élevées, ce qui rend difficile le transfert d'électrons de valence d'un atome à un autre.Considérons une molécule d'ammoniac. L'atome d'azote nécessite plus de trois électrons pour atteindre un octet, et l'atome d'hydrogène a besoin d'un électron à atteindre un duo. Par conséquent, l'atome d'azote se lie avec trois atomes d'hydrogène de telle sorte que l'azote et l'hydrogène atteignent des configurations d'électrons stables.Puisque chacune de ces liaisons partage une paire d'électrons, on l'appelle une seule liaison. La paire d'électrons partagée dans la liaison covalente est appelée une paire de liaison. Les électrons de valence ne participant pas à la liaison sont appelés la paire solitaire ou électrons non liés.Avec 6 électrons de valence, les atomes d'oxygène ont besoin de deux électrons supplémentaires pour atteindre un octet. Par conséquent, deux atomes d'oxygène partagent des paires de deux électrons formant une double liaison. L'azote, d'autre part, partage les trois électrons non appariés dans sa forme diatomique, créant une triple liaison.Les liaisons simples et multiples diffèrent en termes de longueur et de résistance des liaisons. Les obligations triples sont plus courtes que les obligations doubles, qui sont plus courtes que les obligations simples. La force de liaison augmente avec la multiplicité des liaisons.C'est pourquoi il est difficile de briser la triple liaison en azote, ce qui la rend relativement réactif. Le modèle de Lewis aide à prévoir la structure et la stabilité des molécules. Selon le modèle de Lewis, H2O est une molécule stable;parce que l'oxygène et l'hydrogène ont obtenu des configurations d'électrons stables.Si l'oxygène partage des électrons avec un seul atome d'hydrogène, la molécule OH résultante n'est pas stable car l'oxygène n'a que 7 électrons de valence et ne peut atteindre l'octet. Cependant, si un électron supplémentaire est ajouté à l'oxygène pour compléter l'octet, l'ion hydroxyde résultant devient stable avec une charge négative. Les liaisons covalentes sont directionnelles car les électrons partagés relient deux paires spécifiques d'atomes.En revanche, les liaisons ioniques sont non directionnelles et retiennent plusieurs ions dans le réseau. Ainsi, dans un composé covalent, les molécules individuelles sont considérées comme des unités fondamentales.

liaisons covalentes et structures de lewis

Par rapport aux liaisons ioniques, qui résultent du transfert d'électrons entre des atomes métalliques et non métalliques, les liaisons covalentes résultent de l'attraction mutuelle des atomes pour un doublet d'électrons “ commun ”.  
 
 Des liaisons covalentes se forment entre deux atomes lorsqu'ils ont tous deux tendance à attirer des électrons vers eux-mêmes (c.-à-d. lorsque les deux atomes ont des énergies d'ionisation et des affinités électroniques identiques ou assez similaires).  
 
 <h4>Propriétés physiques des composés covalents </h4>
 
 Les composés qui contiennent des liaisons covalentes présentent des propriétés physiques différentes de celles des composés ioniques. Comme l'attraction entre les molécules, qui sont électriquement neutres, est plus faible que celle qui existe entre les ions chargés électriquement, les composés covalents ont généralement des points de fusion et d'ébullition beaucoup plus faibles que les composés ioniques. En fait, de nombreux composés covalents sont des liquides ou des gaz à température ambiante et, à l'état solide, ils sont généralement beaucoup plus mous que les solides ioniques. En outre, alors que les composés ioniques sont de bons conducteurs d'électricité lorsqu'ils sont dissous dans l'eau, la plupart des composés covalents sont insolubles dans l'eau ; comme ils sont électriquement neutres, ils sont de mauvais conducteurs d'électricité dans n'importe quel état. 
 
 <h4>Formation de liaisons covalentes </h4>
 
 Les atomes de non-métaux forment souvent des liaisons covalentes avec d'autres atomes de non-métaux. Par exemple, la molécule d'hydrogène H2 contient une liaison covalente entre ses deux atomes d'hydrogène. Deux atomes d'hydrogène séparés ayant une énergie potentielle particulière s'approchent l'un de l'autre, leurs orbitales de valence (1s) commencent à se chevaucher. Les électrons célibataires sur chaque atome d'hydrogène interagissent alors avec les deux noyaux atomiques, occupant l'espace autour des deux atomes. La forte attraction de chaque électron commun vers les deux noyaux stabilise le système et l'énergie potentielle diminue à mesure que la distance de la liaison diminue. Si les atomes continuent à se rapprocher, les charges positives dans les deux noyaux commencent à se repousser, et l'énergie potentielle augmente. La longueur de la liaison est déterminée par la distance à laquelle l'énergie potentielle la plus faible est atteinte. 
 
 Il est essentiel de se rappeler que l'énergie doit être ajoutée pour rompre les liaisons chimiques (un processus endothermique), alors que la formation de liaisons chimiques libère de l'énergie (un processus exothermique). Dans le cas de H2, la liaison covalente est très forte ; une grande quantité d'énergie, 436 kJ, doit être ajoutée pour rompre les liaisons en une seule mole de molécules d'hydrogène et provoquer la séparation des atomes : 
 
 <img alt="Eq1" src="https://www.jovecore.com/CDNSource/lessons/1059/1059_Equations-01.png" style="height: 40px;" />
 
 Inversement, la même quantité d'énergie est libérée lorsqu'une mole de molécules de H2 se forme à partir de deux moles d'atomes H : 
 
 <img alt="Eq2" src="https://www.jovecore.com/CDNSource/lessons/1059/1059_Equations-02.png" style="height: 40px;" />
 
 <h4>Structures de Lewis</h4>
 
 Les symboles de Lewis peuvent être utilisés pour indiquer la formation de liaisons covalentes, qui sont illustrées dans les structures de Lewis, des dessins qui décrivent la liaison dans les molécules et les ions polyatomiques. Par exemple, lorsque deux atomes de chlore forment une molécule de chlore, ils mettent en commun un doublet d'électrons : 
 
 <img alt="Figure1" src="https://www.jovecore.com/CDNSource/lessons/1059/1059Image_1.svg" style="height: 90px;" />
 
 La structure de Lewis indique que chaque atome de Cl possède trois doublets d'électrons qui ne sont pas utilisées dans la liaison (appelées doublets libres) et un doublet d'électrons commun (inscrit entre les atomes). Un tiret (ou une ligne) est parfois utilisé pour indiquer un doublet d'électrons commun :  Cl&—Cl. 
 
 <ul>
 <li>Un seul doublet d'électrons commun est appelé une liaison simple. Chaque atome de Cl interagit avec huit électrons de valence : les six des doublets libres et les deux de la liaison simple. </li>
 <li>Cependant, une paire d'atomes peut avoir besoin de mettre en commun plus qu'un doublet d'électrons pour acquérir l'octet requis. Une liaison double se forme lorsque deux doublets d'électrons sont communs entre une paire d'atomes, comme entre les atomes de carbone et d'oxygène dans CH2O (formaldéhyde) et entre les deux atomes de carbone dans C2H4 (éthylène).   
 <img alt="Figure2" src="https://www.jovecore.com/CDNSource/lessons/1059/1059Image_2.svg" style="height: 120px;" /></li>
 <li>Une liaison triple se forme lorsque trois doublets d'électrons sont mis en commun par une paire d'atomes, comme dans le monoxyde de carbone (CO) et l'ion cyanure (CN–). 
 <img alt="Figure3" src="https://www.jovecore.com/CDNSource/lessons/1059/1059Image_3.svg" style="height: 90px;" /></li>
 </ul>
 
 Le tableau périodique peut être utilisé pour prédire le nombre d'électrons de valence dans un atome et le nombre de liaisons qui seront formées pour atteindre un octet. Les éléments du groupe 18, tels que l'argon et l'hélium, ont rempli la configuration électronique et participent donc rarement à la liaison chimique. Cependant, les atomes du groupe 17, comme le brome ou l'iode, n'ont besoin que d'un électron pour atteindre l'octet. Par conséquent, les atomes appartenant au groupe 17 peuvent former une seule liaison covalente. Les atomes du groupe 16 ont besoin de 2 électrons pour atteindre un octet; ils peuvent donc former deux liaisons covalentes. De même, le carbone appartenant au groupe 14 a besoin de 4 électrons de plus pour atteindre un octet ; ainsi, le carbone peut former quatre liaisons covalentes.  
 
 Ce texte est adapté de <a href="https://openstax.org/books/chemistry-2e/pages/7-2-covalent-bonding">Openstax, Chimie 2e, Section 7.2 : Liaisons covalentes</a> et <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>

Liaisons covalentes et structures de Lewis

Par rapport aux liaisons ioniques, qui résultent du transfert d’électrons entre les atomes métalliques et non métalliques, les liaisons covalentes résultent de la ...

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shared electrons in covalent molecules

Covalent molecules form when atoms share electrons instead of transferring them. This bonding usually happens between nonmetal atoms with similar tendencies to attract electrons. Because the atoms pull on the electrons in a shared way, the bond is held together by mutual attraction.

Covalent compounds often have physical properties that differ from ionic compounds. The attraction between neutral molecules is weaker than the attraction between charged ions. As a result, many covalent compounds have lower melting and boiling points, and some are liquids or gases at room temperature. When they are solids, they are usually softer than ionic solids. Most covalent compounds are also insoluble in water and are poor conductors of electricity in any state.

Hydrogen gas shows how a covalent bond forms. In H2, two hydrogen atoms approach each other and their 1s valence orbitals overlap. Each electron interacts with both nuclei, which lowers the potential energy and stabilizes the molecule. If the atoms move too close, the positive nuclei repel each other, and the potential energy rises. The bond length is the distance where the potential energy is lowest. Forming a bond releases energy, so it is an exothermic process. Breaking a bond requires energy, so it is endothermic. For H2, 436 kJ must be added to break one mole of molecules, and the same amount is released when one mole forms from hydrogen atoms.

Lewis structures show covalent bonding with Lewis symbols. In chlorine gas, Cl2, the two chlorine atoms share one pair of electrons. Each chlorine atom also has three lone pairs, which are pairs not used in bonding. A shared pair can be shown with a dash, so Cl—Cl represents a single bond. Each chlorine atom then has eight valence electrons around it, including lone pairs and the bonding pair.

Some atoms need more than one shared pair to reach an octet, which is a full outer shell of eight electrons. A double bond forms when two pairs of electrons are shared, as in CH2O and C2H4. A triple bond forms when three pairs are shared, as in CO and CN−. The periodic table helps predict bonding patterns from valence electrons. Group 18 elements such as argon and helium already have filled electron configurations and rarely bond. Group 17 atoms, such as bromine and iodine, usually need one electron and can form one covalent bond. Group 16 atoms need two electrons and can form two covalent bonds. Carbon in group 14 needs four more electrons, so it can form four covalent bonds.

Shared Electrons in Covalent Molecules

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Shared Electrons in Covalent Molecules

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