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

Lewis-Strukturen sind vereinfachte Darstellungen chemischer Bindungen zwischen Atomen. Betrachten Sie Ethen, bei dem zwei Kohlenstoffatome mit jeweils vier Valenzelektronen von vier Wasserstoffatomen mit jeweils einem Valenzelektron umgeben sind. Ethen hat also insgesamt zwölf Valenzelektronen.
Es braucht vierundzwanzig Elektronen, um die Oktette für zwei Kohlenstoffatome und die Duette für vier Wasserstoffatome zu erfüllen. Zwölf Elektronen müssen also zwischen den Atomen aufgeteilt werden. Da nur zwölf Elektronen zur Verfügung stehen, müssen sie alle Bindungselektronen sein, damit alle Atome eine stabile elektronische Konfiguration erreichen können.
Manchmal weisen Atome in mehratomigen Strukturen nicht die Standardanzahl von Valenzelektronen auf. In solchen Fällen ist eine formale Ladung ungleich Null, F, mit dem anomalen Atom verbunden. Die formale Ladung ist gleich der Anzahl der Valenzelektronen auf dem neutralen Atom abzüglich der Anzahl der Bindungen und ungeteilten Elektronen auf diesem Atom.
Im Ammoniumion zum Beispiel hat jedes Wasserstoffatom ein Valenzelektron in seiner neutralen Form und ist durch eine Einfachbindung gebunden, was eine formale Ladung von Null ergibt.
Dadurch hat das Stickstoffatom vier Valenzelektronen statt der üblichen fünf Valenzelektronen. Da ihm ein Elektron fehlt, trägt er eine formale Ladung von eins plus. Somit beträgt die Gesamtladung des Ammoniumions eins plus.
Stabile Konfigurationen minimieren in der Regel die formalen Gebühren. Bei einigen größeren Elementen kann das Oktett so erweitert werden, dass es mehr als acht Elektronen teilt, um eine Konfiguration mit weniger formalen Ladungen zu erreichen.
So kann das Schwefelatom in Schwefeltrioxid sowohl Einfach- als auch Doppelbindungen mit den umgebenden Sauerstoffatomen eingehen. Die Doppelbindungsanordnung wird jedoch bevorzugt, da sie die formale Gesamtladung des Atoms reduziert.
Weitere Ausnahmen von der Oktettregel sind borhaltige Verbindungen, die ein Ziel von sechs statt acht gemeinsamen Elektronen haben, da das Erreichen eines Oktetts eine formale Ladung erfordert.

Adaptiert von <a href="https://openstax.org/books/chemistry-2e/pages/7-3-lewis-symbols-and-structures">Openstax, Chemie 2e, Abschnitt 7.3: Lewis-Symbole und -Strukturen</a> und <a href="https://openstax.org/books/chemistry-2e/pages/7-4-formal-charges-and-resonance">Openstax, Chemie 2e, Abschnitt 7.4: Formale Ladungen und Resonanz</a>.

lewis-strukturen und formalladungen

Lewis-Symbole können verwendet werden, um die Bildung kovalenter Bindungen anzuzeigen, die in Lewis-Strukturen dargestellt werden – Zeichnungen, die die Bindung in Molekülen und mehratomigen Ionen beschreiben. Das Periodensystem kann verwendet werden, um die Anzahl der Valenzelektronen in einem Atom und die Anzahl der Bindungen vorherzusagen, die gebildet werden, um ein Oktett zu erreichen. Elemente der Gruppe 18 wie Argon und Helium haben gefüllte Elektronenkonfigurationen und sind daher selten an chemischen Bindungen beteiligt. Atome der Gruppe 17 wie Brom oder Jod benötigen jedoch nur ein Elektron, um das Oktett zu erreichen. Daher können Atome der Gruppe 17 eine einzelne kovalente Bindung eingehen. Die Atome der Gruppe 16 benötigen zwei Elektronen, um ein Oktett zu erreichen; Daher können sie zwei kovalente Bindungen bilden. Ebenso benötigt Kohlenstoff, der zur Gruppe 14 gehört, vier weitere Elektronen, um ein Oktett zu erreichen; Somit kann Kohlenstoff vier kovalente Bindungen eingehen.

Betrachten Sie die Lewis-Struktur des Chlormoleküls:

<img alt="Figure1" src="/files/ftp_upload/11681/PTfigure_1.svg" style="text-align: center; height: 30px;">

Die Lewis-Struktur weist darauf hin, dass jedes Cl-Atom drei Elektronenpaare hat, die nicht für die Bindung verwendet werden (die sogenannten freien Elektronenpaare) und ein gemeinsames Elektronenpaar (zwischen den Atomen dargestellt). Manchmal wird ein Strich (oder eine Linie) verwendet, um ein gemeinsames Elektronenpaar anzuzeigen: Cl–Cl

Ein einzelnes gemeinsames Elektronenpaar wird als Einfachbindung bezeichnet. Jedes Cl-Atom interagiert mit acht Valenzelektronen: den sechs in den freien Elektronenpaaren und den zwei in der Einfachbindung. Allerdings muss ein Atompaar möglicherweise mehr als ein Elektronenpaar teilen, um das erforderliche Oktett zu erreichen.

Eine Doppelbindung entsteht, wenn zwei Elektronenpaare zwischen einem Atompaar geteilt werden, wie z. B. zwischen den Kohlenstoff- und Sauerstoffatomen in CH2O (Formaldehyd).

<img alt="Figure2" src="/files/ftp_upload/11681/PTfigure_2.svg" style="text-align: center; height: 60px;">

Eine Dreifachbindung entsteht, wenn drei Elektronenpaare von einem Atompaar gemeinsam genutzt werden, wie beispielsweise bei Kohlenmonoxid (CO).

<img alt="Figure3" src="/files/ftp_upload/11681/PTfigure_3.svg" style="text-align: center; height: 20px;">

Die formale Ladung eines Atoms in einem Molekül ist die hypothetische Ladung, die das Atom hätte, wenn die Elektronen in den Bindungen gleichmäßig zwischen den Atomen verteilt wären. Die Formalladung kann berechnet werden, indem man die Summe der Anzahl der nichtbindenden Elektronen und der Anzahl der Bindungen an einem Atom (oder die Hälfte der Anzahl der Bindungselektronen) von der Anzahl der Valenzelektronen des neutralen Atoms abzieht:

Formale Ladung = # Valenzschalenelektronen (freies Atom) − # freie Elektronenpaare − # Bindungen

Die formellen Ladungsberechnungen können noch einmal überprüft werden, indem die Summe der formalen Ladungen für die gesamte Struktur ermittelt wird. Die Summe der formalen Ladungen aller Atome in einem neutralen Molekül muss Null sein; Die Summe der formalen Ladungen in einem Ion sollte der Ladung des Ions entsprechen. Bedenken Sie, dass die für ein Atom berechnete formale Ladung nicht die tatsächliche Ladung des Atoms im Molekül ist. Die Formalladung ist nur ein nützliches Buchhaltungsverfahren; es weist nicht auf das Vorhandensein tatsächlicher Ladungen hin.

Lewis-Strukturen und Formalladungen

Lewis-Symbole können verwendet werden, um die Bildung kovalenter Bindungen anzuzeigen, die in Lewis-Strukturen dargestellt sind – Zeichnungen, die die Bindung in ...

Education

Lehre

JoVE Core

Organic Chemistry

Organische Chemie

Unit 1

Einheit 1

Covalent Bonding and Structure

Kovalente Bindung und Struktur

KEY TERMS AND CONCEPTS

KERNBEGRIFFE UND KONZEPTE

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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CHEM1.PS1.1

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CHEM1.PS1.2

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CHEM1.PS1.3

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Oxidation&ndash;Reduction Reactions

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CHEM1.PS1.4

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CHEM1.PS1.6

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The chelating agent must contain hydrophobic groups and ...

Chromatography: Introduction

Chromatography is a technique used to separate compounds based on differences of partitioning between two phases, the stationary phase and the mobile ...

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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