Lewis diagrams are visual tools used to represent the arrangement of valence electrons in atoms and molecules, highlighting bonding and lone pairs of electrons.
What Are Lewis Diagrams?
Lewis diagrams, also referred to as Lewis dot structures, are a foundational method for visualizing molecular structure in chemistry. They represent the valence electrons—the electrons in the outermost shell of an atom—that are involved in chemical bonding. By showing both the bonding pairs of electrons (shared between atoms) and the lone pairs (non-bonding pairs located on a single atom), Lewis diagrams help explain how atoms connect to form molecules and ions.
These diagrams are built on the octet rule, a principle stating that atoms tend to form bonds in such a way that each atom (except a few exceptions) achieves a full outer shell of eight valence electrons, similar to the configuration of noble gases. The Lewis diagram gives insights into both structure and reactivity, and is essential for predicting molecular shapes, bond types, and reaction mechanisms.
The Lewis structure model is part of the localized electron model, where electrons are considered to be localized between atoms or around individual atoms, rather than moving freely through a structure. This model contrasts with delocalized models such as those seen in metallic bonding or resonance structures.
Components of a Lewis Diagram
Each Lewis diagram includes the following components:
Lone pairs: Also known as nonbonding pairs, these are pairs of valence electrons that are not shared with other atoms. They belong to a single atom and often influence molecular shape and polarity.
Bonding pairs: These are electrons shared between two atoms to form covalent bonds. A single bond involves one pair of shared electrons, a double bond involves two pairs, and a triple bond involves three pairs. Each bond is often represented by a line (–) in diagrams.
These structures provide a simple yet powerful representation of how atoms achieve stable electron configurations through the formation of chemical bonds.
Ionic Lewis Dot Structures
Ionic compounds form when electrons are transferred from one atom to another, creating oppositely charged ions that attract each other. This usually involves a metal donating electrons to become a cation (positively charged ion) and a nonmetal accepting electrons to become an anion (negatively charged ion).
Steps for Drawing Ionic Lewis Dot Structures
Write the empirical formula to identify the elements and the ratio of atoms in the compound.
Determine the number of valence electrons for each atom by consulting their group on the periodic table.
Transfer electrons from the metal to the nonmetal so both achieve full outer electron shells.
Place brackets around each ion to indicate that they are no longer neutral atoms.
Label the charge of each ion outside the brackets, using superscripts like +1 or –1.
Image Courtesy of Wayne Breslyn
Example: Sodium Bromide (NaBr)
Sodium (Na) is in Group 1 and has 1 valence electron.
Bromine (Br) is in Group 17 and has 7 valence electrons.
Sodium donates its one electron to bromine, resulting in:
Na⁺ (a cation with a full shell, now resembling neon)
Br⁻ (an anion with 8 valence electrons, resembling krypton)
The final diagram shows [Na]⁺ and [Br]⁻ enclosed in brackets with the charges indicated outside.
Practice Tip: Try drawing the ionic Lewis structure of magnesium chloride (MgCl₂). Remember, magnesium donates two electrons, and each chlorine accepts one.
Covalent Lewis Dot Structures
In covalent bonding, electrons are shared between atoms instead of transferred. The goal is the same: each atom wants to achieve a stable electron configuration, usually through the octet rule. Covalent bonds are common in organic molecules and biological systems.
Steps for Drawing Covalent Lewis Structures
Count the total number of valence electrons available in the molecule. Add or subtract electrons if the molecule is an ion.
Identify the central atom, typically the one with the lowest electronegativity (except hydrogen, which is never central).
Connect surrounding atoms to the central atom using single bonds (each bond uses two electrons).
Distribute remaining electrons to complete the octets of the outer atoms first.
Place any leftover electrons on the central atom.
If atoms do not have full octets, form double or triple bonds by converting lone pairs into bonding pairs.
Example: O₂ (Dioxygen)
Oxygen is in Group 16 and has 6 valence electrons.
Two oxygen atoms give a total of 12 valence electrons.
Start by connecting the two O atoms with a single bond (2 electrons).
Add lone pairs: 3 pairs per oxygen → 12 electrons total (but 14 used).
Since 14 electrons exceed the available 12, convert one lone pair on each oxygen into a bonding pair, forming a double bond.
This results in both atoms achieving an octet and using exactly 12 electrons.
Example: N₂ (Dinitrogen)
Each nitrogen atom has 5 valence electrons, for a total of 10.
Begin with a single bond and add lone pairs. You will exceed 10 electrons.
Convert lone pairs into bonding pairs until a triple bond forms.
Final structure: one lone pair per nitrogen and a triple bond between them.
Exceptions to the Octet Rule
Although the octet rule is widely applicable, there are several important exceptions to be aware of.
Incomplete Octets
Certain elements are stable with fewer than 8 electrons:
Hydrogen (H): Can only hold 2 electrons.
Beryllium (Be): Stable with 4 electrons.
Boron (B): Usually stable with 6 electrons.
These elements are often found in molecules where the octet rule would not be applicable.
Expanded Octets
Elements in Period 3 and beyond can use d orbitals to accommodate more than 8 electrons.
Examples include phosphorus, sulfur, and xenon.
This is possible when the central atom has an atomic number ≥ 14.
Example: XeF₂ (Xenon Difluoride)
Xenon has 8 valence electrons.
Fluorine has 7 valence electrons × 2 = 14.
Total: 22 valence electrons.
After bonding and filling octets, Xe ends up with 10 electrons, which is permissible due to its expanded shell capability.
Odd-Electron Molecules
Sometimes, molecules have an odd number of total valence electrons, which prevents a perfect pairing.
These are called free radicals.
Example: NO (Nitric Oxide) – has 11 electrons total and cannot satisfy the octet rule for both atoms.
Common Examples of Lewis Structures
Example 1: CS₂ (Carbon Disulfide)
Carbon has 4 valence electrons.
Each sulfur has 6 valence electrons.
Total: 4 + 6×2 = 16 valence electrons.
Carbon is the central atom.
Single bonds between C and each S → 4 electrons used.
Adding full octets puts us at 20 electrons, which is too many.
Convert lone pairs to double bonds.
Final structure: two double bonds between carbon and sulfur.
Example 2: XeF₂
Xenon (Xe) has 8 valence electrons.
Fluorine (F) has 7 each × 2 = 14.
Total: 22 valence electrons.
Single bonds to F use 4 electrons.
Lone pairs on each F take 12 more.
Remaining 6 electrons go to Xe as three lone pairs.
Note: Lone pairs on Xe must be kept in pairs on opposite sides for symmetry and correct representation.
Example 3: NH₄⁺ (Ammonium Ion)
N = 5 valence electrons; H = 1 each × 4 = 4.
Total = 9 valence electrons.
Positive charge (+1) → remove one electron → 8 valence electrons.
Place N in the center and form four single bonds to H.
The entire structure should be placed in brackets with a + sign.
Tips for Drawing Lewis Diagrams
Always double-check electron counts before finalizing the structure.
Hydrogen never forms more than one bond.
Carbon usually forms four bonds and no lone pairs.
Oxygen commonly forms two bonds and two lone pairs.
If central atoms lack a full octet, form double or triple bonds where necessary.
Use brackets and charges when drawing ions.
Avoid placing unpaired electrons on atoms unless the molecule is a known radical.
Practice Questions
Try drawing the following structures using the steps outlined above:
Magnesium chloride (MgCl₂)
Mg = +2; Cl = –1 × 2
Each Cl receives one electron from Mg
Final structure: [Mg]²⁺ and two [Cl]⁻ ions
Image Courtesy of Wayne Breslyn
Nitrogen gas (N₂)
10 valence electrons
Requires a triple bond and one lone pair per nitrogen
Image Courtesy of Clutch Prep
Ammonium ion (NH₄⁺)
8 valence electrons due to the +1 charge
One central N atom with four single bonds to H atoms
No lone pairs, full octet for nitrogen
Image Courtesy of MaketheBrainhappy
Mastering Lewis diagrams is essential for understanding molecular composition, especially in the context of biology where molecules like water, carbon dioxide, and ammonia play vital roles in life processes. Practice frequently and become confident in interpreting and drawing these essential structures.
FAQ
Formal charges are used to evaluate the stability and likelihood of a particular Lewis structure. They help identify which resonance structure or bonding arrangement is most chemically reasonable. The formal charge of an atom is calculated as:
Valence electrons - (Lone pair electrons + 1/2 Bonding electrons)
To determine the most stable structure:
The best Lewis structure usually has formal charges closest to zero.
Negative formal charges should be placed on the most electronegative atoms.
Avoid structures with adjacent atoms carrying like charges.
Structures with minimal charge separation are generally more stable.
When multiple resonance forms exist, the one with the lowest overall formal charges and correct charge placement is preferred.
Using formal charges allows students to distinguish between possible structures and identify the most plausible one based on electronic distribution and atomic properties.
Yes, Lewis diagrams can give insight into molecular polarity when combined with an understanding of molecular geometry. While the diagram alone shows bonding and lone pairs, polarity arises from both electronegativity differences and molecular shape.
Unequal sharing of electrons (polar bonds) occurs between atoms with different electronegativities.
A symmetric arrangement of polar bonds can cancel dipole moments, making the molecule nonpolar.
Asymmetrical shapes (due to lone pairs or uneven atom distribution) usually create a net dipole.
For example, CO2 has polar bonds but is linear and symmetrical, so it's nonpolar.
In contrast, H2O has a bent shape due to lone pairs, resulting in a net dipole and polarity.
Lewis structures help visualize electron domains, which then contribute to determining molecular polarity.
Some molecules cannot be accurately depicted by a single Lewis diagram because electrons are delocalized, meaning they are spread across multiple atoms rather than localized between two atoms.
Resonance structures are multiple Lewis diagrams representing the same molecule, differing only in the arrangement of electrons, not atoms.
The actual structure is a resonance hybrid, a blend of all valid resonance forms.
Molecules like ozone (O3), nitrate ion (NO3−), and benzene (C6H6) require resonance to depict their bonding accurately.
Each resonance structure contributes to the hybrid, stabilizing the molecule.
Curved arrows are often used in advanced chemistry to show electron movement, but in Lewis diagrams, we show all reasonable electron arrangements.
Resonance emphasizes that some bonds are fractional in character—neither purely single nor double—which a single Lewis diagram cannot show on its own.
Elements in Period 3 and beyond can expand their valence shells to accommodate more than 8 electrons because they have available d orbitals.
The octet rule applies well to Period 2 elements (like C, N, O, and F) because their valence shells contain only s and p orbitals.
Period 3 elements (like P, S, Cl) have empty 3d orbitals, which can be used for bonding.
This allows them to form expanded octets and accommodate 10, 12, or more valence electrons.
Examples include:
SF6 (sulfur hexafluoride): sulfur forms six bonds, using 12 valence electrons.
PCl5 (phosphorus pentachloride): phosphorus forms five bonds, with 10 electrons.
In Lewis diagrams, this is shown by adding additional bonding or lone pairs beyond the octet.
Understanding the role of d orbitals is crucial when dealing with larger atoms that break the traditional octet rule.
Polyatomic ions require special handling in Lewis diagrams because they have a net electrical charge due to the gain or loss of electrons.
Determine the total number of valence electrons based on the sum of individual atoms' electrons, then add or subtract electrons based on the charge.
Add electrons for negative charges.
Subtract electrons for positive charges.
Enclose the entire Lewis diagram within square brackets.
Indicate the overall charge as a superscript outside the top right of the brackets (e.g., [NO3]−).
The rest of the process is similar: assign bonding and lone pairs, ensure octets (or exceptions), and form double/triple bonds if needed.
Practice Questions
Draw the Lewis diagram for the polyatomic ion NH4+ and explain how the structure reflects the bonding and electron distribution in the ion.
The Lewis diagram of NH4+ shows a nitrogen atom at the center with four hydrogen atoms surrounding it, each connected by a single covalent bond. Nitrogen starts with five valence electrons and each hydrogen contributes one, totaling nine, but the positive charge indicates the loss of one electron, leaving eight. These eight electrons form four bonding pairs, with no lone pairs on nitrogen. Each hydrogen shares its single electron with nitrogen, completing its duet. The positive charge is shown outside brackets surrounding the structure, reflecting the overall loss of one electron and stable electron configuration through shared bonding.
A student draws the Lewis structure for CO2 as O–C–O. Evaluate the student’s drawing and correct it using proper Lewis structure principles.
The student's drawing of O–C–O is incorrect because it does not provide a full octet for all atoms and misrepresents the bonding in CO2. Each oxygen has six valence electrons, and carbon has four, for a total of 16. The correct Lewis structure involves placing carbon in the center with each oxygen forming a double bond with carbon, using 8 electrons in bonds and 8 as lone pairs. This arrangement ensures all atoms satisfy the octet rule. The revised structure shows carbon with four shared pairs and each oxygen with two lone pairs, accurately reflecting bonding and electron distribution.
