Name: chemical bonds: types and strength
Uploaded: 2023-04-30T18:28:31.000Z
Duration: 2 min 40 s
Channel: JoVE

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Louisiana Student Standards for Science

High School

Chemistry

Matter And Its Interactions

HS-PS1-1

Chemical Bonds

Lorsque les atomes se rapprochent, leurs noyaux se repoussent, tout comme les électrons. En même temps, les électrons de chaque atome sont attirés par le noyau de l'autre et s'organisent de manière à maximiser les forces d'attraction entre les atomes tout en minimisant les forces de répulsion.
Si cela se traduit par une force d’attraction nette, l’énergie potentielle du système est réduite. Cette interaction s’appelle une liaison chimique.
Cependant, toutes les combinaisons d’atomes ne peuvent pas former des liaisons chimiques stables. Le comportement chimique d’un atome est principalement déterminé par ce qui forme des liaisons et le type de liaisons dont il s’agit. Comprendre la liaison entre les atomes qui composent un composé aide à comprendre et à prédire leur comportement moléculaire.
L’énergie de liaison est l’énergie nécessaire pour rompre une liaison en une mole d’un composé gazeux. Cela dépend du type d’atomes liés et du nombre d’électrons impliqués. La force d’adhérence augmente avec l’énergie d’adhérence.
Le type d’atomes liés et le nombre d’électrons impliqués influencent également la longueur de la liaison, qui est la distance moyenne entre les noyaux de deux atomes liés et est inversement proportionnelle à la multiplicité et à la force de la liaison.
Les interactions de liaison les plus courantes se situent sur un spectre allant des liaisons ioniques, qui impliquent des attractions électrostatiques entre des ions de charge opposée, aux liaisons covalentes, qui partagent des électrons de valence entre les atomes.
La liaison ionique est la principale interaction trouvée dans les composés ioniques. Les cations et les anions constitutifs sont maintenus ensemble par des forces d’attraction électrostatiques. Souvent, la formation d’ions est précédée par le transfert d’électrons de valence entre les atomes participants.
L’atome qui perd des électrons devient le cation, tandis que l’atome qui accepte les électrons devient l’anion. Le transfert d’électrons donne généralement à chaque ion la configuration de gaz noble la plus proche, c’est-à-dire la configuration atteinte avec le plus petit changement dans le nombre d’électrons, ce qui rend les ions énergétiquement plus stables.
Dans la liaison covalente, les électrons partagés interagissent avec les noyaux des atomes de liaison et réduisent l’énergie potentielle. Les liaisons doubles et triples partagent respectivement deux et trois paires d’électrons entre les atomes. Comme cela augmente le nombre d’électrons soumis à des forces d’attraction supplémentaires de l’autre noyau, la force de liaison augmente avec la multiplicité des liaisons.

liaisons chimiques

Les atomes participent à la formation de liaisons chimiques  pour acquérir une configuration électronique de couche de valence complétée similaire à celle du gaz noble le plus proche en numéro atomique. Les liaisons ioniques, covalentes et métalliques sont quelques-uns des types importants de liaisons chimiques. L'énergie de liaison et la longueur de liaison déterminent la force d'une liaison chimique.

<h4>Types de liaisons chimiques</h4>

Une liaison ionique est formée en raison de l'attraction électrostatique entre des cations et des anions. Souvent, les ions sont formés par le transfert d'électrons d'un atome participant à l'autre. Cependant, ces liaisons n'ont pas de directionnalité définie car la force d'attraction électrostatique est uniformément répartie dans l'espace tridimensionnel.

Une liaison covalente est une liaison chimique formée par le partage de paires d'électrons entre des atomes adjacents. La paire d'électrons partagée est appelée paire liante. Les liaisons covalentes sont directionnelles par nature.

Une liaison métallique est formée entre deux atomes métalliques. La liaison métallique est décrite par le modèle de la « Mer d'électrons ». Sur la base des faibles énergies d'ionisation des métaux, le modèle indique que les atomes métalliques perdent facilement leurs électrons de valence et deviennent des cations. Ces électrons de valence créent un pool d'électrons délocalisés entourant les cations sur tout le métal.

<h4>Énergies de liaison et longueur de liaison</h4>

La force d'une liaison covalente est mesurée par l'énergie nécessaire pour la rompre, c'est-à-dire l'énergie nécessaire pour séparer les atomes liés. Séparer une paire d'atomes liés nécessite de l'énergie. Plus une liaison est forte, plus l'énergie nécessaire pour la rompre est grande.

L'énergie nécessaire pour rompre une liaison covalente spécifique dans une mole de molécules gazeuses est appelée énergie de liaison ou énergie de dissociation de liaison. L'énergie de liaison pour une molécule diatomique est définie comme le changement d'enthalpie standard pour la réaction endothermique. Les molécules avec trois atomes ou plus ont deux liaisons ou plus. La somme de toutes les énergies de liaison dans une telle molécule est égale au changement d'enthalpie standard pour la réaction endothermique qui rompt toutes les liaisons dans la molécule.

La force d'une liaison entre deux atomes augmente à mesure que le nombre de paires d'électrons dans la liaison augmente. Généralement, plus le nombre de liaisons entre deux atomes est élevé, plus la longueur de liaison est courte et plus la force de liaison est grande. Ainsi, les liaisons triples sont plus fortes et plus courtes que les liaisons doubles entre les mêmes deux atomes ; de même, les liaisons doubles sont plus fortes et plus courtes que les liaisons simples entre les mêmes deux atomes. Lorsqu'un atome se lie à divers atomes dans un groupe, la force de liaison diminue généralement à mesure que nous descendons dans le groupe.

Ce texte est adapté de <a href="https://openstax.org/books/chemistry-2e/pages/7-1-ionic-bonding">Openstax, Chemistry 2e, Section 7.1: Ionic Bonding</a>, <a href="https://openstax.org/books/chemistry-2e/pages/7-2-covalent-bonding">Openstax, Section 7.2: Covalent Bonding</a>, <a href="https://openstax.org/books/chemistry-2e/pages/10-5-the-solid-state-of-matter">Section 10.5: The Solid State of Matter</a>, and Section<a href="https://openstax.org/books/chemistry-2e/pages/7-5-strengths-of-ionic-and-covalent-bonds"> 7.5. Bond Strength: Covalent Bonds</a>.

Liaisons chimiques


Les atomes participent à la formation d’une liaison chimique pour acquérir une configuration électronique complète de couche de valence similaire à celle du gaz noble le plus proche ...

Education

Enseignement

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Organic Chemistry

Chimie organique

Unit 1

Unité 1

Covalent Bonding and Structure

Liaison covalente et structure

KEY TERMS AND CONCEPTS

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chemical bonds: types and strength

Chemical bonds help atoms reach a full valence shell like the nearest noble gas. The main bond types are ionic, covalent, and metallic. Bond energy and bond length also help explain how strong a bond is.

An ionic bond forms from electrostatic attraction between cations and anions. In many cases, one atom transfers electrons to another atom, and the ions form from that transfer. Ionic bonds have no defined direction because the attraction spreads evenly through three-dimensional space.

A covalent bond forms when adjacent atoms share electron pairs. The shared pair is called the bonding pair. Covalent bonds are directional, so the atoms are linked in a specific orientation.

A metallic bond forms between two metal atoms. The electron sea model describes this type of bonding. In this model, metal atoms lose valence electrons easily and become cations, while the delocalized electrons form a pool around the cations across the metal.

The strength of a covalent bond is measured by the energy needed to break it and separate the bonded atoms. Stronger bonds require more energy. For one mole of gaseous molecules, this value is called bond energy, or bond dissociation energy. For a diatomic molecule, it is the standard enthalpy change for the endothermic reaction that breaks the bond.

Molecules with three or more atoms have more than one bond. In those cases, the total bond energy is the sum of all bond energies needed to break every bond in the molecule. Bond strength also increases as the number of electron pairs in the bond increases. Triple bonds are stronger and shorter than double bonds between the same atoms, and double bonds are stronger and shorter than single bonds. When one atom bonds to different atoms in a group, bond strength usually decreases as you move down the group.

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chemical bonds

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Chemical Bonds: Types and Strength

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