Name: lewis structures, bonds, and formal charge
Uploaded: 2025-05-22T09:05:32.000Z
Duration: 2 min 19 s
Channel: JoVE

K12 Standards

Tennessee Academic Standards

High School

Chemistry

Matter and Its Interactions

CHEM1.PS1.12

Lewis Structures and Formal Charges

Les structures de Lewis sont des représentations simplifiées des liaisons chimiques entre les atomes. Considérons l’éthène, où deux atomes de carbone de quatre électrons de valence chacun sont entourés de quatre atomes d’hydrogène ayant chacun un électron de valence. Par conséquent, l’éthylène a un total de douze électrons de valence.
Il faut vingt-quatre électrons pour satisfaire les octets de deux atomes de carbone et les duos de quatre atomes d’hydrogène. Ainsi, douze électrons doivent être partagés entre les atomes. Avec seulement douze électrons disponibles, ils doivent tous être des électrons de liaison pour permettre à tous les atomes d’atteindre une configuration électronique stable.
Parfois, les atomes dans les structures polyatomiques ne présentent pas le nombre standard d’électrons de valence. Dans de tels cas, une charge formelle non nulle, F, est associée à l’atome anormal. La charge formelle est égale au nombre d’électrons de valence sur l’atome neutre moins le nombre de liaisons et d’électrons non partagés sur cet atome.
Par exemple, dans l’ion ammonium, chaque atome d’hydrogène a un électron de valence sous sa forme neutre et est lié par une seule liaison, ce qui donne une charge formelle de zéro.
Cela laisse l’atome d’azote avec quatre électrons de valence au lieu de ses cinq électrons de valence habituels. Puisqu’il lui manque un électron, il porte une charge formelle de un plus. Ainsi, la charge globale de l’ion ammonium est de un plus.
Les configurations stables minimisent généralement les frais formels. Pour certains éléments plus grands, l’octet peut être élargi pour partager plus de huit électrons afin d’atteindre une configuration avec moins de charges formelles.
Par exemple, l’atome de soufre dans le trioxyde de soufre peut former des liaisons simples et doubles avec les atomes d’oxygène environnants. Cependant, l’arrangement à double liaison est préféré, car il réduit la charge formelle globale sur l’atome.
D’autres exceptions à la règle de l’octet incluent les composés contenant du bore, qui ont une cible de six électrons partagés, au lieu de huit, car l’obtention d’un octet nécessite une charge formelle.

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> et <a href="https://openstax.org/books/chemistry-2e/pages/7-4-formal-charges-and-resonance">Openstax, Chimie 2e, Section 7.4 : Charges formelles et résonance</a>.

structures de lewis et accusations formelles

Les symboles de Lewis peuvent être utilisés pour indiquer la formation de liaisons covalentes, qui sont représentées dans les structures de Lewis, des dessins qui décrivent la liaison dans les molécules et les ions polyatomiques. 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 des configurations électroniques remplies et participent donc rarement aux liaisons chimiques. Cependant, les atomes du groupe 17, comme le brome ou l’iode, n’ont besoin que d’un seul électron pour atteindre l’octet. Ainsi, les atomes appartenant au groupe 17 peuvent former une seule liaison covalente. Les atomes du groupe 16 ont besoin de deux électrons pour atteindre un octet ; par conséquent, ils peuvent former deux liaisons covalentes. De même, le carbone, qui appartient au groupe 14, a besoin de quatre électrons supplémentaires pour atteindre un octet ; ainsi, le carbone peut former quatre liaisons covalentes.

Considérons la structure de Lewis de la molécule de chlore :

<img alt="Figure1" src="/files/ftp_upload/11681/PTfigure_1.svg" style="text-align: center; height: 30px;">

La structure de Lewis indique que chaque atome de Cl possède trois paires d'électrons qui ne sont pas utilisées dans la liaison (appelées paires isolées) et une paire d'électrons partagée (écrite entre les atomes). Un tiret (ou une ligne) est parfois utilisé pour indiquer une paire d'électrons partagée : Cl–Cl

Une seule paire d’électrons partagée est appelée une simple liaison. Chaque atome Cl interagit avec huit électrons de valence : les six dans les paires libres et les deux dans la liaison simple. Cependant, une paire d’atomes peut devoir partager plus d’une paire d’électrons pour obtenir l’octet requis.

Une double liaison se forme lorsque deux paires d’électrons sont partagées entre une paire d’atomes, comme entre les atomes de carbone et d’oxygène dans CH2O (formaldéhyde).

<img alt="Figure2" src="/files/ftp_upload/11681/PTfigure_2.svg" style="text-align: center; height: 60px;">

Une triple liaison se forme lorsque trois paires d'électrons sont partagées par une paire d'atomes, comme dans le monoxyde de carbone (CO).

<img alt="Figure3" src="/files/ftp_upload/11681/PTfigure_3.svg" style="text-align: center; height: 20px;">

La charge formelle d’un atome dans une molécule est la charge hypothétique que l’atome aurait si les électrons des liaisons étaient uniformément répartis entre les atomes. La charge formelle peut être calculée en soustrayant la somme du nombre d'électrons non liants et du nombre de liaisons sur un atome (ou la moitié du nombre d'électrons liants) du nombre d'électrons de valence de l'atome neutre :

Charge formelle = # électrons de couche de valence (atome libre) − # électrons isolés − # liaisons

Les calculs des frais formels peuvent être revérifiés en déterminant la somme des frais formels pour l’ensemble de la structure. La somme des charges formelles de tous les atomes d’une molécule neutre doit être nulle ; la somme des charges formelles d'un ion doit être égale à la charge de l'ion. N'oubliez pas que la charge formelle calculée pour un atome n'est pas la charge réelle de l'atome dans la molécule. La taxe formelle n'est qu'une procédure comptable utile ; cela n’indique pas la présence de frais réels.

Structures de Lewis et accusations formelles

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 ...

Education

Enseignement

JoVE Core

Organic Chemistry

Chimie organique

Unit 1

Unité 1

Covalent Bonding and Structure

Liaison covalente et structure

KEY TERMS AND CONCEPTS

TERMES ET CONCEPTS CLÉS

lewis structures, bonds, and formal charge

Lewis structures show how atoms share electrons in covalent bonds and in polyatomic ions. Lewis symbols help draw these structures, and the periodic table helps predict how many valence electrons an atom has and how many bonds it can form to reach an octet.

Group 18 elements, such as argon and helium, already have filled electron configurations. They rarely take part in chemical bonding. Atoms in group 17, such as bromine and iodine, need one more electron to reach an octet, so they can form one covalent bond. Atoms in group 16 need two electrons, so they can form two covalent bonds. Carbon is in group 14 and needs four more electrons, so it can form four covalent bonds.

The chlorine molecule is a simple example of a Lewis structure. Each Cl atom has three lone pairs, which are electron pairs not used in bonding, and one shared pair between the atoms. A dash or line can show this shared pair, so the bond can be written as Cl–Cl. A single shared pair is a single bond, and each chlorine atom reaches eight valence electrons when its lone pairs and bond are counted.

Some atoms must share more than one pair of electrons to complete an octet. A double bond forms when two pairs of electrons are shared, as in formaldehyde, CH2O, between carbon and oxygen. A triple bond forms when three electron pairs are shared, as in carbon monoxide, CO. These bond types help show how atoms connect in a molecule.

Formal charge is a bookkeeping tool for Lewis structures. It is the hypothetical charge an atom would have if bonding electrons were shared evenly. To calculate it, subtract the number of lone pair electrons and the number of bonds from the number of valence shell electrons in the free atom. The formal charges in a neutral molecule must add to zero, and the formal charges in an ion must add to the ion’s charge. Even so, formal charge is not the same as actual charge on an atom.

Lewis Structures, Bonds, and Formal Charge

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Atomic Orbitals and Quantum Numbers

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Lewis structures show how atoms share electrons in covalent bonds and in polyatomic ions. Lewis symbols help draw these structures, and the periodic table helps predict how many valence electrons an atom has and how many bonds it can form to reach an octet. Group 18 elements, such as argon and helium, already have filled electron configurations. They rarely take part in chemical bonding. Atoms in group 17, such as bromine and iodine, need one more electron to reach an octet, so they can form...

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Predicting Shape and Molecular Polarity

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resonance and hybrid structures

Resonance in Nitrite and Carbonate Ions

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intermolecular forces and physical properties

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Functional Groups and Organic Molecules

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key terms and concepts

Learn how Lewis structures show bonding, lone pairs, single, double, and triple bonds, and how to calculate formal charge in molecules and ions.

Lewis Structures, Bonds, and Formal Charge | JoVE K12

organic chemistry

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CHEM1.PS1.1

Chemical Equations

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The Atomic Theory of Matter

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Demonstrate that atoms, and therefore mass, are conserved during a chemical reaction by balancing chemical reactions. EXPLANATION: Standard 7.PS1.4 explores the idea that chemical reactions are rearrangements of atoms, and students undertake the process of balancing chemical reactions with that standard in seventh grade. The introduction of molar masses in high school makes an explicit connection between the number of atoms, as previously investigated, and the actual masses of the atoms, permitting the collection of more concrete evidence for the conservation of mass in chemical reactions.

CHEM1.PS1.2

Reaction Yield

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Limiting Reactant

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chemical-equilibria-systematic-approach-to-equilibrium-calculations

Chemical Equilibria: Systematic Approach to Equilibrium Calculations

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Perform stoichiometric calculations involving the following relationships: mole- mole; mass-mass; mole-mass; mole- particle; and mass- particle. Show a qualitative understanding of the phenomenon of percent yield, limiting, and excess reagents in a chemical reaction through pictorial and conceptual examples. (states of matter liquid and solid; excluding volume of gasses). EXPLANATION: This standard builds on the idea that balancing chemical reactions provides evidence for conservation of mass and that the behavior of atoms follows predictable patterns. Students now have the opportunity to utilize this understanding as they perform and evaluate chemical reactions. (The concepts of limiting reagent and percent yield are presented; however, students are not expected to determine these values algebraically. Use of molar volume in stoichiometric processes is beyond the scope of this standard.)

CHEM1.PS1.3

Predicting Reaction Outcomes

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Introduction to Chemical Reactions

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Chemical Reactions in Aqueous Solutions

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Precipitation Reactions

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Oxidation-Reduction Reactions

Oxidation&ndash;Reduction Reactions

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Oxidation Numbers

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Acids, Bases and Neutralization Reactions

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CHEM1.PS1.4

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Kinetic Molecular Theory and Gas Laws Explain Properties of Gas Molecules

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Basic Postulates of Kinetic Molecular Theory: Particle Size, Energy, and Collision

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Ideal Gas Law

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Conduct investigations to explore and characterize the behavior of gasses (pressure, volume, and temperature), develop models to represent this behavior, and construct arguments to explain this behavior. Evaluate the relationship (qualitative and quantitatively) at STP between pressure and volume (Boyle's Law), temperature and volume (Charles' Law), temperature and pressure (Gay-Lussac Law) and moles and volume (Avogadro's Law), and evaluate and explain these relationships with respect to kinetic molecular theory. Be able to understand, establish, and predict the relationships between volume, temperature, and pressure using combined gas law both qualitatively and quantitatively. EXPLANATION: When exploring the behavior of gasses, it is important to consider experimental design. Experiments used to show the relationships between these sets of variables should include one independent and one dependent variable. Other variables should be held constant. Pressure should serve as the dependent variable because it cannot be manipulated directly. Avogadro's Law can be used to relate the amount of gas in two different containers in lieu of knowing actual number of gas particles in an individual container. Individual demonstrations can be performed to explore each of the different gas laws.

CHEM1.PS1.5

Ideal Gas Equation

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Chemical Stoichiometry and Gases: Using Ideal Gas Law to Determine Moles

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Use the ideal gas law, PV = nRT, to algebraically evaluate the relationship among the number of moles, volume, pressure, and temperature for ideal gases. EXPLANATION: If students have collected data during investigations for CHEM1.PS1.5, these data can then be used to linearize. With pressure assigned as the dependent variable and plotted against (moles) (temperature)/volume, the slope of the graph will be the gas constant, yielding the ideal gas law. After developing the mathematical model, this model can be used to make predictions for other systems.

CHEM1.PS1.6

Separation of Mixtures via Precipitation

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Filtration

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Centrifugation

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Precipitation and Co-precipitation

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Dialysis

Dialysis separates desired analyte ions or molecules from unwanted molecules, or &#39;interferents&#39;, by allowing the sample solution to diffuse ...

Analyze solutions to identify solutes and solvents, quantitatively analyze concentrations (molarity, percent composition, and ppm), and perform separation methods such as evaporation, distillation, and/or chromatography and show conceptual understanding of distillation. Construct argument to justify the use of certain separation methods under different conditions. EXPLANATION: In seventh grade, students developed an understanding that chemical and physical properties of a sample of matter could be used to identify a substance. Building on this understanding, students now explore the use of these physical and chemical properties to separate mixtures. Analysis of the substances is performed to aid in further discussion and facilitate further quantitative analysis.

CHEM1.PS1.7

Lewis Structures, Bonds, and Formal Charge

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