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

As estruturas de Lewis são representações simplificadas de ligações químicas entre átomos. Considere o eteno, onde dois átomos de carbono com quatro elétrons de valência cada são cercados por quatro átomos de hidrogênio com um elétron de valência cada. Portanto, o eteno tem um total de doze elétrons de valência.
São necessários vinte e quatro elétrons para satisfazer os octetos de dois átomos de carbono e os duetos de quatro átomos de hidrogênio. Assim, doze elétrons devem ser compartilhados entre os átomos. Com apenas doze elétrons disponíveis, todos eles devem estar ligando elétrons para permitir que todos os átomos atinjam uma configuração eletrônica estável.
Às vezes, os átomos em estruturas poliatômicas não exibem o número padrão de elétrons de valência. Nesses casos, uma carga formal diferente de zero, F, está associada ao átomo anômalo. A carga formal é igual ao número de elétrons de valência no átomo neutro menos o número de ligações e elétrons não compartilhados nesse átomo.
Por exemplo, no íon amônio, cada átomo de hidrogênio tem um elétron de valência em sua forma neutra e é ligado por uma ligação simples, produzindo uma carga formal de zero.
Isso deixa o átomo de nitrogênio com quatro elétrons de valência em vez de seus cinco elétrons de valência usuais. Como está faltando um elétron, ele carrega uma carga formal de um mais. Assim, a carga geral do íon amônio é de um a mais.
Configurações estáveis geralmente minimizam as cobranças formais. Para alguns elementos maiores, o octeto pode ser expandido para compartilhar mais de oito elétrons para atingir uma configuração com menos cargas formais.
Por exemplo, o átomo de enxofre no trióxido de enxofre pode formar ligações simples e duplas com os átomos de oxigênio circundantes. No entanto, o arranjo de ligação dupla é preferido, pois reduz a carga formal geral no átomo.
Outras exceções à regra do octeto incluem compostos contendo boro, que têm um alvo de seis elétrons compartilhados, em vez de oito, porque atingir um octeto requer uma carga formal.

Adaptado de <a href="https://openstax.org/books/chemistry-2e/pages/7-3-lewis-symbols-and-structures">Openstax, Química 2e, Seção 7.3: Símbolos e Estruturas de Lewis</a> e <a href="https://openstax.org/books/chemistry-2e/pages/7-4-formal-charges-and-resonance">Openstax, Química 2e, Seção 7.4: Encargos Formais e Ressonância</a>.

estruturas de lewis e cargas formais

Os símbolos de Lewis podem ser usados ​​para indicar a formação de ligações covalentes, que são mostradas nas estruturas de Lewis – desenhos que descrevem a ligação em moléculas e íons poliatômicos. A tabela periódica pode ser usada para prever o número de elétrons de valência em um átomo e o número de ligações que serão formadas para atingir um octeto. Elementos do grupo 18, como argônio e hélio, possuem configurações eletrônicas preenchidas e, portanto, raramente participam de ligações químicas. No entanto, os átomos do grupo 17, como o bromo ou o iodo, precisam de apenas um elétron para atingir o octeto. Consequentemente, os átomos pertencentes ao grupo 17 podem formar uma única ligação covalente. Os átomos do grupo 16 precisam de dois elétrons para atingir um octeto; portanto, eles podem formar duas ligações covalentes. Da mesma forma, o carbono, que pertence ao grupo 14, precisa de mais quatro elétrons para atingir um octeto; assim, o carbono pode formar quatro ligações covalentes.

Considere a estrutura de Lewis da molécula de cloro:

<img alt="Figure1" src="/files/ftp_upload/11681/PTfigure_1.svg" style="text-align: center; height: 30px;">

A estrutura de Lewis indica que cada átomo de Cl possui três pares de elétrons que não são usados ​​na ligação (chamados de pares solitários) e um par de elétrons compartilhado (escrito entre os átomos). Às vezes, um traço (ou linha) é usado para indicar um par compartilhado de elétrons: Cl–Cl

Um único par compartilhado de elétrons é chamado de ligação simples. Cada átomo de Cl interage com oito elétrons de valência: os seis nos pares solitários e os dois na ligação simples. No entanto, um par de átomos pode precisar compartilhar mais de um par de elétrons para atingir o octeto necessário.

Uma ligação dupla se forma quando dois pares de elétrons são compartilhados entre um par de átomos, como entre os átomos de carbono e oxigênio no CH2O (formaldeído).

<img alt="Figure2" src="/files/ftp_upload/11681/PTfigure_2.svg" style="text-align: center; height: 60px;">

Uma ligação tripla se forma quando três pares de elétrons são compartilhados por um par de átomos, como no monóxido de carbono (CO).

<img alt="Figure3" src="/files/ftp_upload/11681/PTfigure_3.svg" style="text-align: center; height: 20px;">

A carga formal de um átomo em uma molécula é a carga hipotética que o átomo teria se os elétrons nas ligações estivessem distribuídos uniformemente entre os átomos. A carga formal pode ser calculada subtraindo a soma do número de elétrons não-ligantes e o número de ligações em um átomo (ou metade do número de elétrons ligantes) do número de elétrons de valência do átomo neutro:

Carga formal = # elétrons da camada de valência (átomo livre) - # par de elétrons solitários - # ligações

Os cálculos de carga formal podem ser verificados novamente determinando a soma das cargas formais para toda a estrutura. A soma das cargas formais de todos os átomos de uma molécula neutra deve ser zero; a soma das cargas formais de um íon deve ser igual à carga do íon. Lembre-se de que a carga formal calculada para um átomo não é a carga real do átomo na molécula. A carga formal é apenas um procedimento contábil útil; não indica a presença de cargas reais.

Estruturas de Lewis e Cargas Formais

Os símbolos de Lewis podem ser usados para indicar a formação de ligações covalentes, que são mostradas em estruturas de Lewis - desenhos que descrevem a ligação em ...

Education

JoVE Core

Organic Chemistry

Química Orgânica

Unit 1

Unidade 1

Covalent Bonding and Structure

Ligação Covalente e Estrutura

KEY TERMS AND CONCEPTS

TERMOS-CHAVE E CONCEITOS

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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How Electrons Fill Atomic Orbitals

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chemical bonds

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...

polar covalent bonds

Electronegativity in Polar Covalent Bonds

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lewis structures and formal charges

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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Molecular Shapes from Electron Pairs

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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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valence bond theory and hybridized orbitals

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Covalent Bond Overlap and Orbital Hybridization

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Bond Order and Molecular Orbitals

Molecular orbital theory explains how electrons are arranged in molecules. It treats electron regions in a molecule much like atomic orbitals in an atom. A molecular orbital is the region of space where a valence electron is most likely to be found. Molecular orbitals form through the linear combination of atomic orbitals, or LCAO. When atomic wave functions combine in phase, they create regions with high electron density. When they combine out of phase, they create nodes, or regions with no...

intermolecular forces and physical properties

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introduction to functional groups

Functional Groups and Organic Molecules

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Advanced Functional Groups in Organic Chemistry

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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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Determining the Mass Percent Composition in an Aqueous Solution

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Solutions and Concentrations

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Stoichiometry, Product Yield, and Limiting Reactants

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Understand and be prepared to use values specific to chemical processes: the mole, molar mass, molarity, and percent composition. EXPLANATION: During the early and mid-1800s, a tremendous amount of early research poured into building an understanding of the way that atoms combine. This work required analysis of the percent composition of different substances. Joseph Proust's work with tin oxides led him to publish his law of definite proportions noting that a given pure substance will have a constant ratio for the combination of elements making up the substance. Building on Proust's work, John Dalton published his law of multiple proportions and atomic theory noting that multiple patterns exist for the way that atoms combine, yet the combining atoms are invariable. Meanwhile Amedeo Avogadro hypothesizes that given equal pressures and temperatures, two equal volumes of gas will contain equal numbers of atoms. Acknowledging that the exact number of atoms was unknown in these gaseous volumes, the Latin word "mole" was used. "Mole" translates to "lump" recognizing the ambiguous quantity of gas. A number of attempts have been undertaken historically to quantify the mole. Robert Millikan's oil drop experiment revealed the change of a single electron, which could then be divided into Faraday's Constant for the charge of a mole of electrons. Dividing the charge of a mole of electrons by the charge of a single electron yields the number of electrons in a mole, $6.022\times10^{\wedge}23$. Students should be familiar with how to translate between the given quantities. Student's first use atomism/Dalton's atomic theory in third grade noting that matter is made up of particles too small to be seen. This the first time where students will undertake the process of counting these invisible particles. Electrolysis of water provides a tangible experience to demonstrate the constituent parts of water, though the definite proportions of water may not be easily seen due to the solubility of the gasses produced. The concept of relative mass can be explored by comparing small groups of like items. For example, relative masses of beans can be compared with pinto, white, and black beans. If bags are created with equal numbers of like beans, students can develop a concept of why the smallest bean will have a relative mass of 1, analogous to hydrogen.

CHEM1.PS1.1

Chemical Equations

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The Atomic Theory of Matter

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Chemical Reactions

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Balancing Redox Equations

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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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Synthesis and Decomposition Reactions

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Types of Chemical Reactions: Exchange and Reversible

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Use the reactants in a chemical reaction to predict the products and identify reaction classes (synthesis, decomposition, combustion, single replacement, double replacement). EXPLANATION: Chemical reactions are rearrangements of atoms that follow predicable patterns. There are patterns both at the macroscopic level in the behavior of some of the reaction classes, as well as patterns in the rearrangements of the atoms underlying the reaction. Students should be able to predict the products of the reactions, which also requires the ability to recognize the general patterns for each type of reaction.

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

Derivation of the Ideal Gas Law
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Kinetic Theory of an Ideal Gas

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Applications of the Ideal Gas Law: Molar Mass, Density, and Volume

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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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Real Gases: Effects of Intermolecular Forces and Molecular Volume Deriving Van der Waals Equation

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

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pV-Diagrams

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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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Distillation: Vapor&ndash;Liquid Equilibria

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Filtration

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Centrifugation

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Sublimation

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Recrystallization: Solid&ndash;Solution Equilibria

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Precipitation and Co-precipitation

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Electrodeposition

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Extraction: Effects of pH

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Extraction: Advanced Methods

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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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