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 structures are simplified representations of chemical bonds between atoms. Consider ethene, where two carbon atoms with four valence electrons each are surrounded by four hydrogen atoms having one valence electron each. Hence, ethene has a total of twelve valence electrons.
It takes twenty-four electrons to satisfy the octets for two carbon atoms and the duets for four hydrogen atoms. Thus, twelve electrons must be shared between the atoms. With only twelve electrons available, they must all be bonding electrons to allow all atoms to reach a stable electronic configuration.
Sometimes, atoms in polyatomic structures do not exhibit the standard number of valence electrons. In such cases, a nonzero formal charge, F, is associated with the anomalous atom. The formal charge is equal to the number of valence electrons on the neutral atom minus the number of bonds and unshared electrons on that atom.
For example, in the ammonium ion, each hydrogen atom has one valence electron in its neutral form and is bonded by a single bond, yielding a formal charge of zero.
This leaves the nitrogen atom with four valence electrons instead of its usual five valence electrons. Since it is missing one electron, it bears a formal charge of one-plus. Thus, the overall charge on the ammonium ion is one-plus.
Stable configurations generally minimize formal charges. For some larger elements, the octet may be expanded to share more than eight electrons to reach a configuration with fewer formal charges.
For example, the sulfur atom in sulfur trioxide can form both single and double bonds with the surrounding oxygen atoms. However, the double bond arrangement is preferred, as it reduces the overall formal charge on the atom.
Other exceptions to the octet rule include boron-bearing compounds, which have a target of six shared electrons, rather than eight, because achieving an octet requires a formal charge.

lewis structures and formal charges

Lewis symbols can be used to indicate the formation of covalent bonds, which are shown in Lewis structures&#8212;drawings that describe the bonding in molecules and polyatomic ions. The periodic table can be used to predict the number of valence electrons in an atom and the number of bonds that will be formed to reach an octet. Group 18 elements, such as argon and helium, have filled electron configurations and thus rarely participate in chemical bonding. However, atoms from group 17, such as bromine or iodine, need only one electron to reach the octet. Hence, atoms belonging to group 17 can form a single covalent bond. The atoms of group 16 need two electrons to reach an octet; hence, they can form two covalent bonds. Similarly, carbon, which belongs to group 14, needs four more electrons to reach an octet; thus, carbon can form four covalent bonds.
Consider the Lewis structure of the chlorine molecule:
<img alt="Chlorine molecule Lewis structure diagram, showing electron sharing and covalent bond formation." src="/files/ftp_upload/11681/PTfigure_1.svg" style="text-align: center; height: 30px;" data-altupdatedat="1754904545000" data-alt="Chlorine molecule Lewis structure, Cl&#8322;, with lone pairs diagram for chemical bonding." />
The Lewis structure indicates that each Cl atom has three pairs of electrons that are not used in bonding (called lone pairs) and one shared pair of electrons (written between the atoms). A dash (or line) is sometimes used to indicate a shared pair of electrons: Cl&#8211;Cl
A single shared pair of electrons is called a single bond. Each Cl atom interacts with eight valence electrons: the six in the lone pairs and the two in the single bond. However, a pair of atoms may need to share more than one pair of electrons in order to achieve the requisite octet.
A double bond forms when two pairs of electrons are shared between a pair of atoms, as between the carbon and oxygen atoms in CH2O (formaldehyde).
<img alt="Molecular structure, formaldehyde chemical formula CH2O, diagram for organic chemistry study." src="/files/ftp_upload/11681/PTfigure_2.svg" style="text-align: center; height: 60px;" data-altupdatedat="1754904545000" data-alt="Formaldehyde molecular structure, chemical diagram, covalent bonds, organic compound formula." />
A triple bond forms when three electron pairs are shared by a pair of atoms, as in carbon monoxide (CO).
<img alt="Carbon monoxide (CO) molecule diagram showing triple bond structure with lone pairs." src="/files/ftp_upload/11681/PTfigure_3.svg" style="text-align: center; height: 20px;" data-altupdatedat="1754904545000" data-alt="CO2 Lewis structure, chemical bonding, linear molecule, valence electrons, molecular diagram." />
The formal charge of an atom in a molecule is the hypothetical charge the atom would have if the electrons in the bonds are evenly distributed between the atoms. The formal charge can be calculated by subtracting the sum of the number of nonbonding electrons and the number of bonds on an atom (or half the number of bonding electrons) from the number of valence electrons of the neutral atom:
Formal charge = # valence shell electrons (free atom) &#8722; # lone pair electrons &#8722; # bonds
The formal charge calculations can be double-checked by determining the sum of the formal charges for the whole structure. The sum of the formal charges of all atoms in a neutral molecule must be zero; the sum of the formal charges in an ion should equal the charge of the ion. Remember that the formal charge calculated for an atom is not the actual charge of the atom in the molecule. The formal charge is only a useful bookkeeping procedure; it does not indicate the presence of actual charges.

Lewis symbols can be used to indicate the formation of covalent bonds, which are shown in Lewis structures&#8212;drawings that describe the bonding in ...

Education

JoVE Core

Organic Chemistry

Unit 1

Covalent Bonding and Structure

KEY TERMS AND CONCEPTS

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

what is organic chemistry?

Carbon Compounds in Everyday Life

Organic chemistry focuses on carbon compounds called organic compounds. These compounds may come from living things or be made by chemists. Carbon is the main element in them, and it bonds with other carbon atoms as well as hydrogen, oxygen, nitrogen, and sulfur. Carbon can form up to four strong bonds. That ability lets carbon atoms build chains and rings of many sizes, shapes, and levels of complexity. Not every carbon-containing substance is classified as organic. Carbonates, cyanides, and...

electronic structure of atoms

Atomic Orbitals and Quantum Numbers

Atomic orbitals and quantum numbers describe how electrons are arranged around the nucleus. An atom has a small, positively charged nucleus with electrons surrounding it. The nucleus contains protons and neutrons, and the quantum mechanical model gives a better picture of this structure than older models. Protons carry a +1 charge and have a mass of 1.0073 amu. Neutrons have no charge and have a mass of 1.0087 amu. Electrons carry a −1 charge and have a mass of 0.00055 amu. The number of...

electron configurations

How Electrons Fill Atomic Orbitals

Electron filling in atoms follows a set order based on orbital energy. The Aufbau principle says each new electron enters the lowest-energy subshell available. The Pauli exclusion principle says no two electrons can share the same set of four quantum numbers. Hund’s rule of maximum multiplicity says electrons stay unpaired in equal-energy orbitals when possible. The order of filling follows the relative energies of the subshells. A common sequence is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on.

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

Polar covalent bonds form when atoms share electrons unevenly. This happens in covalent bonds between atoms that have similar, but not identical, tendencies to attract electrons. Nonmetal atoms often make these bonds with other nonmetals. Hydrogen gas, H2, is a simple example of a covalent bond between two hydrogen atoms. A covalent bond begins when two atoms move close enough for their valence orbitals to overlap. In hydrogen, the 1s orbitals overlap and the single electrons on each atom...

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...

vsepr theory

Molecular Shapes from Electron Pairs

Molecular shape can be predicted from the way electron pairs arrange around a central atom. Valence shell electron-pair repulsion theory, or VSEPR theory, uses a Lewis structure to count bonding pairs and lone electron pairs. The main idea is simple: electron pairs spread out as far as possible to reduce repulsion. In a molecule, the valence electrons around the central atom can be bonding pairs or lone pairs. Bonding pairs are found mainly between bonded atoms. Lone pairs stay on one atom and...

molecular geometry and dipole moments

Predicting Shape and Molecular Polarity

Molecular shape and molecular polarity can be predicted with VSEPR theory and dipole moments. Start by writing the Lewis structure of the molecule or polyatomic ion. Then count the electron groups around the central atom. Lone pairs and bonds both count as electron groups, and a single, double, or triple bond each counts as one region of electron density. Use that count to identify the electron-pair geometry. Next, use the number of lone pairs to determine the molecular structure. If more than...

resonance and hybrid structures

Resonance in Nitrite and Carbonate Ions

Resonance explains why some Lewis structures do not match the measured properties of a molecule or ion. When two or more Lewis structures have the same atom arrangement, the real electron distribution is an average of the structures shown. The nitrite anion, NO2−, is one example. It can be drawn in two ways, with the N–O and N=O bonds in different places. If nitrite had one single bond and one double bond, the two bond lengths should be different, because a double bond is shorter and stronger...

valence bond theory and hybridized orbitals

valence-bond-theory-and-hybridized-orbitals

Covalent Bond Overlap and Orbital Hybridization

Covalent bonds form when orbitals on two atoms overlap and their single electrons pair up. The bond is stronger when the overlap is greater. Maximum overlap happens when the orbitals meet along a direct line between the two nuclei. A sigma bond, or σ bond, is the kind of covalent bond found in a single bond in a Lewis structure. In a σ bond, electron density is concentrated along the internuclear axis. A pi bond, or π bond, forms when two p orbitals overlap side by side. In a π bond, the...

mo theory and covalent bonding

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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How Intermolecular Forces Affect Matter

Intermolecular forces are the attractions between molecules, and they help explain how matter behaves. These forces influence melting points, boiling points, and whether substances mix with one another. Water has a much higher boiling point than hexane, which shows that water has stronger intermolecular forces. Three main kinds of intermolecular interactions are covered here. They include ion-dipole forces, dipole-dipole interactions, and van der Waals forces, which include London dispersion...

solubility

Why Some Substances Dissolve in Water

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introduction to functional groups

Functional Groups and Organic Molecules

Functional groups are groups of atoms with specific chemical properties that appear within organic molecules. They are sometimes shown as "R." These groups sit along the carbon backbone of macromolecules, which can form chains or rings of carbon atoms. Functional groups help functionalize a compound. That means they can give the molecule different physical and chemical properties. In organic chemistry, this change in group structure can strongly affect how a compound behaves. Common...

overview of advanced functional groups

Advanced Functional Groups in Organic Chemistry

Advanced functional groups are atom clusters in organic molecules that give those molecules specific chemical properties. They are sometimes shown as "R" in formulas. By changing the atoms attached to a carbon skeleton, a functional group can make a compound behave differently in physical and chemical reactions. This topic covers several major classes of organic compounds with advanced functional groups. Aromatic hydrocarbons include benzene. Phenol is another example and is also called...

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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Formula Mass and Mole Concepts of Compounds

Formula Mass of Covalent Compounds

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Expressing Solution Concentration

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Determining the Mass Percent Composition in an Aqueous Solution

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Solutions and Concentrations

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Reaction Stoichiometry

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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

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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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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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Ideal Gas Law

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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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CHEM1.PS1.6

Separation of Mixtures via Precipitation

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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

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Acid/Base Strengths and Dissociation Constants

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Identify acids and bases as a special class of compounds with a specific set of properties. EXPLANATION: The concept of pH is first introduced in seventh grade as a chemical property of matter. Students should be introduced to multiple explanations of acid and base behavior to permit classification of common substances (e.g. baking soda, ammonia, carbon dioxide) as acids or bases. To further develop discussions of solubility, qualitative discussions of the differences between strong and weak acids are appropriate. (Calculations of pH or ion concentrations are beyond the scope of this standard. Discussions of pOH are beyond the scope of this standard. Conjugate pairs are beyond the scope of this standard.)

CHEM1.PS1.8

Lewis Structures, Bonds, and Formal Charge

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