Getting Started
Covalent compounds are formed when atoms share electrons to achieve stable electron configurations, creating molecules. At the atomic scale, understanding how this sharing occurs is crucial for predicting a molecule's structure and properties. The core problem is visualizing this arrangement of shared and unshared electrons, a challenge addressed by a simple yet powerful two-dimensional model called the Lewis diagram.
What You Should Be Able to Do
By the end of this section, you should be ableto:
Calculate the total number of valence electrons available for bonding in a molecule or polyatomic ion.
Construct a valid Lewis diagram for a given chemical formula that follows the octet rule.
Correctly place single, double, and triple bonds as needed to satisfy the electron requirements of atoms.
Identify and draw plausible Lewis diagrams for common exceptions to the octet rule.
Key Concepts & Analysis
Drawing a Lewis diagram is a systematic process. By treating the chemical formula and valence electron counts as inputs, we can follow a set of steps to produce a valid structural representation as the output. The octet rule serves as the primary control, guiding the arrangement of electrons.
Inputs & Preconditions
Chemical Formula: The identity and number of each atom in the molecule or polyatomic ion (e.g., H₂O, NH₄⁺).
Valence Electrons: The electrons in the outermost energy level of an atom, which are involved in chemical bonding. The number of valence electrons for a main-group element corresponds to its group number in the periodic table (e.g., Carbon in Group 14 has 4; Oxygen in Group 16 has 6; Fluorine in Group 17 has 7).
Total Charge (for ions): For polyatomic ions, the overall charge dictates an adjustment to the total electron count.
Key Steps / Mechanism
Constructing a Lewis diagram follows a reliable, step-by-step procedure. Let's use carbon dioxide (CO₂) as a running example.
Sum the Valence Electrons: Add the valence electrons from all atoms in the formula. For polyatomic ions, add one electron for each unit of negative charge and subtract one for each unit of positive charge.
- Example (CO₂): Carbon (Group 14) has 4 valence electrons. Oxygen (Group 16) has 6. Total = 4 + 2(6) = 16 valence electrons.
Identify the Central Atom & Draw a Skeleton: The central atom is typically the least electronegative element that is not hydrogen. Connect the other atoms (terminal atoms) to the central atom with single bonds. Each single bond uses two electrons.
- Example (CO₂): Carbon is less electronegative than oxygen, so it is the central atom. The skeleton is O—C—O. This uses 4 electrons (2 bonds × 2 e⁻/bond), leaving 16 - 4 = 12 electrons.
Distribute Remaining Electrons to Terminal Atoms: Place the remaining electrons as lone pairs (pairs of nonbonding electrons) on the terminal atoms until each has a complete octet (8 electrons). Hydrogen is an exception, requiring only a duet (2 electrons).
- Example (CO₂): Distribute the 12 remaining electrons to the two oxygen atoms. Each oxygen gets 6 electrons (3 lone pairs), satisfying their octets.
Place Any Remaining Electrons on the Central Atom: If any electrons are still left after filling the terminal atoms' octets, place them on the central atom.
- Example (CO₂): All 12 electrons have been used. There are none left for the central carbon atom.
Check the Central Atom's Octet & Form Multiple Bonds: If the central atom does not have a complete octet, create multiple bonds (double or triple). Convert a lone pair from a terminal atom into a shared bonding pair with the central atom. Repeat until the central atom satisfies the octet rule, the principle that atoms are most stable when they have eight valence electrons.
- Example (CO₂): The carbon atom only has 4 electrons (from the two single bonds). To satisfy its octet, one lone pair from each oxygen atom is moved to form a double bond with carbon. The final structure is O=C=O, with two lone pairs on each oxygen. Now, C has 8 electrons, and each O has 8 electrons.
Outputs & Effects
A Complete Lewis Diagram: The final output is a 2D model showing the arrangement of all valence electrons as either bonding pairs (lines) or lone pairs (dots).
Bond Order: The diagram reveals the number of bonds between atoms (single, double, or triple), which is related to bond length and strength.
Foundation for 3D Geometry: While not a 3D model itself, the Lewis diagram is the essential first step for predicting the three-dimensional shape of a molecule.
Controls & Limiting Factors
The Octet Rule: This is the primary guiding principle. Most atoms (especially C, N, O, F) must end up with eight valence electrons.
Total Valence Electron Count: This number is absolute. The final diagram must contain exactly this many electrons, no more and no less.
Exceptions to the Octet Rule: Certain elements are known exceptions that override the standard octet rule.
Incomplete Octet: Boron (B) and Beryllium (Be) are often stable with fewer than eight electrons (e.g., BF₃, where B has 6).
Expanded Octet: Elements in the third period and below (e.g., P, S, Cl) can accommodate more than eight electrons in their valence shell because they have accessible d-orbitals (e.g., SF₆, where S has 12).
Key Models & Representations
The process of drawing a Lewis diagram can be visualized as a decision-making flowchart.
| Step | Action | Question / Checkpoint |
|---|---|---|
| 1. Foundation | Sum all valence electrons from the chemical formula. Adjust for any ionic charge. | Is the total count correct? |
| 2. Skeleton | Identify the central atom (least electronegative) and connect all terminal atoms with single bonds. | How many electrons were used? Subtract this from the total. |
| 3. Distribution | Distribute remaining electrons as lone pairs, starting with the terminal atoms to satisfy their octets. | Do all terminal atoms (except H) have 8 electrons? |
| 4. Final Check | Place any leftover electrons on the central atom. Check if the central atom has a complete octet. | Does the central atom have at least 8 electrons? (Or is it a known exception?) |
| 5. Refinement | If the central atom's octet is incomplete, move lone pairs from terminal atoms to form double or triple bonds. | Is the central atom's octet now satisfied? Is the total electron count still correct? |
Key Terms, Quantities, & Concepts
Lewis Diagram: A two-dimensional model representing the sharing of valence electrons in a molecule, showing bonding electrons as lines and nonbonding electrons as dots.
Valence Electrons: The electrons in the outermost shell of an atom that participate in chemical bonding.
Octet Rule: The tendency of main-group atoms to form bonds in ways that give them eight valence electrons, achieving a stable configuration similar to that of a noble gas.
Bonding Pair: A pair of electrons shared between two atoms, forming a covalent bond. A single line in a Lewis diagram represents one bonding pair.
Lone Pair (Nonbonding Pair): A pair of valence electrons that is not involved in bonding and belongs to a single atom.
Single Bond: A covalent bond in which one pair of electrons is shared between two atoms.
Double Bond: A covalent bond in which two pairs of electrons are shared between two atoms.
Triple Bond: A covalent bond in which three pairs of electrons are shared between two atoms.
Expanded Octet: A condition where a central atom in a molecule has more than eight valence electrons, possible for elements in the third period and below.
Skill Snapshots
Causation:
Cause: A molecule has a negative charge (it is an anion). Effect: Additional electrons must be added to the total valence electron count before drawing the structure.
Cause: The central atom in an initial Lewis sketch has only six electrons. Effect: A lone pair from an adjacent terminal atom must be converted into a bonding pair to form a double bond, satisfying the central atom's octet.
Cause: The central atom is sulfur, an element in the third period. Effect: It is capable of forming an expanded octet, accommodating more than eight valence electrons, as seen in SF₆.
Comparison:
A bonding pair is shared between two atomic nuclei, holding the molecule together, while a lone pair is localized on a single atom and influences the molecule's shape and reactivity.
A double bond consists of four shared electrons and is stronger and shorter than a single bond, which consists of two shared electrons.
Molecules like CH₄ strictly adhere to the octet rule, whereas molecules like PCl₅ are exceptions, featuring an expanded octet on the central phosphorus atom.
CCOT (Continuity and Change Over Time):
Baseline: A set of individual atoms, each with its own count of valence electrons.
Change 1: The formation of a skeletal structure connects the atoms with single bonds, fundamentally changing their relationship from separate entities to a connected molecule.
Change 2: The conversion of lone pairs into multiple bonds alters the bond order between atoms, strengthening the connection to satisfy the octet rule.
Continuity: The total number of valence electrons calculated in the first step remains constant throughout the entire drawing process; electrons are merely rearranged, never created or destroyed.
Common Misconceptions & Clarifications
Misconception: Every atom in a Lewis diagram must have exactly eight electrons.
- Clarification: This is false. Hydrogen is stable with two electrons (a duet). Boron is often stable with six, and elements from the third period down can have expanded octets (10, 12, or even 14 electrons). The octet rule is a powerful guideline, not an unbreakable law.
Misconception: The Lewis diagram shows the molecule's 3D shape.
- Clarification: A Lewis diagram is a 2D representation of electron accounting and connectivity. It does not accurately depict bond angles or the true three-dimensional geometry of the molecule. That is the purpose of VSEPR theory, which uses the Lewis diagram as its starting point.
Misconception: Forgetting to adjust the electron count for ions.
- Clarification: The charge of a polyatomic ion is critical. For an anion like sulfate (SO₄²⁻), you must add two electrons to the total valence count. For a cation like ammonium (NH₄⁺), you must subtract one electron. Failing to do this will always result in an incorrect diagram.
Misconception: Any atom can be the central atom.
- Clarification: Hydrogen can only form one bond, so it is always a terminal atom. Among other atoms, the least electronegative one is typically the central atom because it is best able to share its electrons with multiple other atoms.
One-Paragraph Summary
Lewis diagrams are fundamental chemical models used to represent the arrangement of valence electrons in covalent molecules and polyatomic ions. The construction of a valid diagram is a systematic process that begins with summing the total available valence electrons and arranging atoms with the least electronegative element at the center. By distributing electrons first as single bonds and then as lone pairs, the primary goal is to satisfy the octet rule for as many atoms as possible. When necessary, lone pairs are converted into double or triple bonds to ensure the central atom achieves a stable octet. While this model is a powerful 2D tool for understanding connectivity and bond order, it is crucial to recognize its limitations and the existence of known exceptions, such as incomplete and expanded octets.