Lewis dot structures visually represent the bonding between atoms in a molecule and show lone pairs of electrons. These diagrams offer a simplified view of valence shell electrons, helping to understand electron arrangement and chemical bonding.
Fundamental Concepts
Valence electrons are the electrons in an atom’s outermost shell. Being furthest from the nucleus, they are least tightly held and participate in chemical bond formation. An element’s valence electron count dictates its chemical properties. For main-group elements, this count is generally the element’s group number on the periodic table. For instance, Group 1 has one valence electron, while Group 14 has four.
The octet rule states that main-group elements tend to achieve eight electrons in their outermost valence shell when bonding. This configuration, similar to noble gases, generally leads to stability. In covalent bonds, shared electrons count towards the octet of both participating atoms. Hydrogen is an exception, following the duet rule by achieving stability with two electrons in its valence shell.
Drawing Lewis Structures Step-by-Step
To construct a Lewis structure, first determine the total number of valence electrons for all atoms. For a neutral molecule, sum the valence electrons of each atom. For a polyatomic ion, add electrons for each negative charge or subtract for each positive charge.
Next, arrange the atoms into a skeletal structure. The central atom is typically the one forming the most bonds or the least electronegative; hydrogen and fluorine are almost always terminal, bonding to only one other atom. Draw single bonds between the central atom and all surrounding atoms. Each single bond represents a shared pair of two valence electrons, which are then subtracted from the total count.
Distribute remaining valence electrons as lone pairs, first on outer atoms to satisfy their octets (or duets for hydrogen), then on the central atom. If the central atom still lacks a complete octet, convert lone pairs from adjacent atoms into double or triple bonds. This allows all atoms (except hydrogen) to satisfy the octet rule and achieve stability.
Applying the Rules: Examples
For water (H₂O), oxygen contributes six valence electrons, and each hydrogen contributes one, totaling eight. Oxygen is the central atom, bonded to two hydrogen atoms. Drawing single bonds uses four electrons, leaving four. These four electrons are placed as two lone pairs on the oxygen, completing its octet and the hydrogen duets.
For carbon dioxide (CO₂), carbon supplies four valence electrons, and each of the two oxygen atoms provides six, totaling sixteen. Carbon is the central atom, with single bonds to each oxygen. This uses four electrons, leaving twelve. Distributing these twelve electrons as lone pairs on the oxygen atoms completes their octets, leaving carbon with only four electrons. To satisfy carbon’s octet, one lone pair from each oxygen forms a double bond with carbon, resulting in two double bonds, with no lone pairs on carbon and two on each oxygen.
For nitrogen (N₂), each nitrogen atom has five valence electrons, totaling ten. With two atoms, they are directly bonded. A single bond uses two electrons, leaving eight electrons to distribute. Placing lone pairs to satisfy octets would require more than the available electrons. Therefore, to achieve octets for both nitrogen atoms, three shared pairs form a triple bond. This leaves one lone pair on each nitrogen, satisfying their octets with the total of ten electrons.
Advanced Considerations
For polyatomic ions, adjust the total valence electron count by adding electrons for each negative charge or subtracting for each positive charge. This ensures the overall charge of the ion is correctly reflected.
Exceptions to the octet rule include incomplete octets, expanded octets, and molecules with an odd number of electrons. Atoms like boron and beryllium can form stable compounds with fewer than eight valence electrons (e.g., six or four). Elements from the third period and beyond, such as sulfur or phosphorus, can accommodate more than eight valence electrons due to available d-orbitals.
When a single Lewis structure cannot fully represent electron distribution, resonance structures are used. These are multiple Lewis structures showing the same atom arrangement but different electron placements (pi bonds or lone pairs). The actual molecule is a hybrid of these structures, with electrons delocalized over several atoms, contributing to stability.