Lewis Dot Structures Practice: Step-by-Step Guide with Examples

Lewis dot structures, also known as electron dot diagrams, are two-dimensional representations of molecules that show the arrangement of valence electrons around atoms. Developed by American chemist Gilbert N. Lewis in 1916, these diagrams remain one of the most fundamental tools in chemistry for visualizing chemical bonding and molecular geometry. Whether you are taking introductory chemistry, organic chemistry, or preparing for standardized exams, lewis dot structures practice is essential for building a strong foundation in chemical reasoning.

A lewis dot structure uses dots to represent valence electrons and lines to represent covalent bonds between atoms. Each line corresponds to a shared pair of electrons, while unshared pairs (lone pairs) are shown as pairs of dots on individual atoms. The goal of a lewis structure is to show how atoms in a molecule share or transfer electrons to achieve a stable electron configuration, typically an octet for main-group elements.

Lewis structures provide critical information that extends beyond simple bonding patterns. They allow chemists to predict molecular geometry using VSEPR theory, determine formal charges on individual atoms, identify resonance structures, and assess whether a molecule is polar or nonpolar. For students preparing for the AP Chemistry exam, the MCAT, or general chemistry finals, fluency in drawing and interpreting lewis structures is non-negotiable.

The power of the electron dot diagram lies in its simplicity. With just a periodic table and knowledge of a few rules, you can represent the bonding in molecules ranging from simple diatomic gases like H2 and O2 to complex polyatomic ions like sulfate and phosphate. In the following sections, we will cover the rules for drawing lewis structures, walk through step-by-step examples, and provide lewis dot structures practice problems to solidify your skills.

Before diving into lewis structure practice problems, it is important to understand the fundamental rules that govern how these diagrams are constructed. These rules are based on the behavior of valence electrons and the tendency of atoms to achieve noble gas electron configurations.

The first rule is the octet rule, which states that most main-group elements tend to form bonds until they are surrounded by eight valence electrons. Hydrogen is a notable exception, as it only requires two electrons to fill its 1s orbital, following the duet rule. Elements in the third period and beyond can sometimes accommodate more than eight electrons due to the availability of empty d orbitals, a phenomenon known as an expanded octet.

The second rule involves counting total valence electrons. To draw any lewis dot structure, you must first determine the total number of valence electrons contributed by all atoms in the molecule or ion. For neutral molecules, simply add the valence electrons of each atom. For polyatomic anions, add one electron for each negative charge. For polyatomic cations, subtract one electron for each positive charge.

The third rule is to identify the central atom. In most molecules, the least electronegative atom (excluding hydrogen) serves as the central atom. Hydrogen and fluorine are always terminal atoms because they form only one bond. Carbon, nitrogen, sulfur, and phosphorus commonly serve as central atoms.

The fourth rule is to distribute electrons to satisfy the octet rule. After placing bonding pairs between the central atom and each surrounding atom, distribute remaining electrons as lone pairs, starting with the terminal atoms. If the central atom does not have an octet after distributing all lone pairs, convert lone pairs on adjacent atoms into additional bonding pairs to form double or triple bonds.

The fifth rule is to check formal charges. The formal charge on each atom is calculated as: valence electrons minus lone pair electrons minus half the bonding electrons. The best lewis structure is the one that minimizes formal charges and places any negative formal charges on the most electronegative atoms. Mastering these rules is the foundation of effective lewis structures practice.

Mastering lewis dot structures requires a systematic approach. By following a consistent step-by-step method, you can accurately draw the lewis dot structure for virtually any molecule or polyatomic ion. This guide walks through the process using water (H2O) and carbon dioxide (CO2) as examples.

Step one is to count the total number of valence electrons. For H2O, each hydrogen contributes one valence electron (2 total) and oxygen contributes six, giving a total of eight valence electrons. For CO2, carbon contributes four and each oxygen contributes six, totaling sixteen valence electrons.

Step two is to determine the central atom and draw a skeleton structure. In H2O, oxygen is the central atom with two hydrogen atoms bonded to it. In CO2, carbon is the central atom with an oxygen on each side. Draw single bonds between the central atom and each terminal atom. Each single bond uses two electrons.

Step three is to distribute remaining electrons as lone pairs. After placing the skeleton bonds, subtract the electrons used in bonding from the total. For H2O, four electrons are used in two O-H bonds, leaving four electrons to distribute as two lone pairs on oxygen. Each hydrogen has its duet satisfied. For CO2, four electrons are used in two C-O bonds, leaving twelve electrons. Place three lone pairs on each oxygen (six electrons each), using all twelve remaining electrons.

Step four is to verify octets. In H2O, oxygen has two bonding pairs and two lone pairs, totaling eight electrons. Both hydrogens have two electrons. The structure is complete. In CO2, each oxygen has an octet, but carbon has only four electrons from two single bonds. To resolve this, convert one lone pair from each oxygen into an additional bonding pair, forming two double bonds. The final electron dot diagram shows O=C=O, with two lone pairs on each oxygen.

Step five is to calculate formal charges and verify the structure is optimal. In the final CO2 structure, all atoms have formal charges of zero, confirming this is the best lewis structure. This systematic approach scales to more complex molecules and is the cornerstone of effective lewis dot structures practice.

The best way to develop fluency with lewis structures is through deliberate practice. Below are several lewis structure practice problems of increasing difficulty, along with detailed solutions. Work through each problem on your own before checking the answer.

Problem 1: Draw the lewis dot structure for ammonia (NH3). Solution: Nitrogen has five valence electrons and each hydrogen has one, totaling eight electrons. Place nitrogen as the central atom with three N-H single bonds (six electrons used). The remaining two electrons form one lone pair on nitrogen. The final structure shows nitrogen with three bonding pairs and one lone pair, giving it a tetrahedral electron geometry and a trigonal pyramidal molecular shape.

Problem 2: Draw the lewis dot structure for the carbonate ion (CO3^2-). Solution: Carbon has four valence electrons, each oxygen has six, and two additional electrons come from the 2- charge, totaling twenty-four electrons. Place carbon at the center bonded to three oxygens. After distributing lone pairs, carbon lacks an octet with single bonds only. Form one double bond between carbon and one oxygen. This yields three resonance structures where the double bond rotates among the three carbon-oxygen positions. Each resonance structure has formal charges of zero on carbon, zero on the double-bonded oxygen, and -1 on each single-bonded oxygen.

Problem 3: Draw the electron dot diagram for sulfur hexafluoride (SF6). Solution: Sulfur has six valence electrons and each fluorine has seven, totaling forty-eight electrons. Sulfur is the central atom bonded to six fluorine atoms. Each fluorine gets three lone pairs after forming an S-F single bond. Sulfur uses twelve electrons in bonding and has zero lone pairs, giving it twelve electrons total — an expanded octet. This is permitted because sulfur is in the third period and has accessible d orbitals.

Problem 4: Draw the lewis structure for ozone (O3). Solution: Three oxygen atoms contribute six valence electrons each, totaling eighteen. The central oxygen bonds to two terminal oxygens. After distributing lone pairs, the central oxygen has only six electrons with single bonds. Convert one lone pair from a terminal oxygen into a double bond. This gives two resonance structures, and the actual ozone molecule is a resonance hybrid with bond order of 1.5 between each oxygen pair.

These lewis dot structures practice problems cover single bonds, double bonds, resonance, polyatomic ions, and expanded octets. Practice each type repeatedly until you can complete them without referencing the steps.

While the octet rule is a powerful guideline for drawing lewis dot structures, several important exceptions exist that students must understand to handle all types of lewis structure practice problems. These exceptions fall into three categories: incomplete octets, odd-electron species, and expanded octets.

Incomplete octets occur when an atom is stable with fewer than eight electrons. The most common examples involve boron and beryllium. Boron trifluoride (BF3) is a classic case: boron has only three valence electrons and forms three single bonds with fluorine, resulting in six electrons around boron. Although resonance structures can be drawn with a B=F double bond that gives boron an octet, experimental evidence shows that the single-bonded structure with an incomplete octet is the dominant contributor. Aluminum behaves similarly in compounds like AlCl3.

Odd-electron species, also called free radicals, contain an unpaired electron and cannot satisfy the octet rule for all atoms. Nitrogen dioxide (NO2) is a well-known example with a total of seventeen valence electrons. Since seventeen is an odd number, at least one atom must have fewer than eight electrons. In the lewis structure of NO2, nitrogen has a lone electron rather than a lone pair. Free radicals are generally highly reactive due to their unpaired electron and participate in chain reactions of biological and industrial significance.

Expanded octets are observed in elements from the third period and beyond, including sulfur, phosphorus, chlorine, and xenon. These atoms have empty d orbitals in their valence shell that can participate in bonding, allowing them to accommodate ten, twelve, or even more electrons. Phosphorus pentachloride (PCl5) places ten electrons around phosphorus, while sulfur hexafluoride (SF6) places twelve. When drawing lewis structures for these molecules, it is important to recognize that expanded octets are not possible for second-period elements such as carbon, nitrogen, and oxygen.

Understanding these exceptions is critical for accurate lewis dot structures practice. When you encounter a molecule that seems impossible to draw using standard rules, check whether the central atom can support an expanded octet or whether you are dealing with a radical species. These exceptions are frequently tested on standardized chemistry exams and appear regularly in lewis structure practice sessions.

Even students who understand the theory behind lewis structures often make preventable errors during lewis dot structures practice. Identifying and correcting these common mistakes before exam day can significantly improve your accuracy and speed.

The most frequent mistake is miscounting valence electrons. Students sometimes forget to add extra electrons for negative charges on polyatomic ions or subtract electrons for positive charges. Always double-check your total before proceeding. A related error is confusing valence electrons with total electrons — remember that only outer-shell electrons are represented in an electron dot diagram.

Another common error is placing too many or too few bonds around hydrogen. Hydrogen always forms exactly one bond and never has lone pairs. If you find hydrogen with two bonds or a lone pair in your lewis structure, something has gone wrong. Similarly, students sometimes place hydrogen as the central atom, which is incorrect — hydrogen can only be a terminal atom.

Failing to minimize formal charges is a subtle but important mistake. When multiple valid lewis structures are possible, the best structure is the one with the smallest formal charges on all atoms. Additionally, negative formal charges should reside on the most electronegative atoms. Students who skip the formal charge calculation often end up with a suboptimal structure that does not accurately represent the molecule.

Ignoring resonance is another pitfall. Many molecules and ions, including ozone, carbonate, nitrate, and benzene, cannot be accurately represented by a single lewis dot structure. When two or more equivalent arrangements of electrons are possible, you must draw all resonance structures and recognize that the actual molecule is a resonance hybrid. Failing to identify resonance leads to incorrect predictions about bond lengths and strengths.

Finally, students sometimes apply the expanded octet to second-period elements. Carbon, nitrogen, oxygen, and fluorine never exceed eight electrons because they lack accessible d orbitals. If your lewis structure shows ten electrons around carbon, revisit your electron distribution. With consistent practice and attention to these common pitfalls, your lewis structures skills will become second nature. LectureScribe's AI-powered practice modules offer targeted feedback on these exact error patterns, helping you build accuracy through repeated lewis dot structures practice.