Lewis structures, also known as electron dot structures, are visual tools used to represent the bonding within molecules. They illustrate the placement of electrons, showing which are shared as chemical bonds and which remain as non-bonding lone pairs. This two-dimensional representation is fundamental for predicting a molecule’s overall shape and understanding its reactivity. Introduced by Gilbert N. Lewis in 1916, these diagrams provide a simple framework for visualizing covalent bonding.
Fundamental Principles of Electron Accounting
The foundation of drawing Lewis structures rests on understanding valence electrons, which are the electrons in an atom’s outermost shell. These electrons are the only ones involved in chemical bonding and form the total pool available for the structure. Every atom must contribute its valence electrons to the total count before bonds or lone pairs are placed.
The octet rule states that atoms generally seek a stable configuration of eight valence electrons. Atoms achieve this stable state by sharing electrons in covalent bonds. The most common exception is Hydrogen, which only requires two electrons to complete its outer shell, following the duplet rule.
The central atom must be identified before drawing the structure. It is typically the least electronegative atom in the molecule, excluding Hydrogen, which is always a terminal atom. Often, the atom that appears only once in the chemical formula, such as Carbon in \(\text{CO}_2\), occupies the central position.
Step-by-Step Procedure for Neutral Molecules
The process begins by calculating the total number of valence electrons contributed by all atoms in the neutral molecule. For instance, in carbon dioxide (\(\text{CO}_2\)), Carbon contributes four valence electrons, and the two Oxygen atoms contribute six each, totaling \(4 + (2 \times 6) = 16\) valence electrons. This total represents the electron budget for the final Lewis structure.
The next step is to determine the skeletal structure by placing the central atom (Carbon in \(\text{CO}_2\)) and arranging the other atoms around it. Single bonds are then drawn between the central atom and each surrounding atom, using two electrons per bond. For \(\text{CO}_2\), two single bonds consume four electrons, leaving 12 electrons remaining.
The remaining electrons are distributed to the terminal atoms first to satisfy their octets. In \(\text{CO}_2\), six electrons (three lone pairs) are added to each Oxygen atom, consuming all 12 remaining electrons.
Any leftover electrons are then placed on the central atom as lone pairs. The structure is checked to ensure every atom has an octet; in \(\text{CO}_2\), the terminal Oxygen atoms are satisfied, but the central Carbon atom only has four electrons.
If the central atom lacks a complete octet, lone pairs from the terminal atoms are moved to form multiple bonds, such as double or triple bonds. For \(\text{CO}_2\), one lone pair from each Oxygen is shifted to form a second bond with the Carbon atom, resulting in two double bonds. This final structure shows the Carbon and Oxygen atoms each surrounded by eight electrons, completing the representation.
Formal Charge and Drawing Polyatomic Ions
The procedure for drawing polyatomic ions, such as the sulfate ion (\(\text{SO}_4^{2-}\)), requires modifying the initial electron count. If the ion carries a negative charge, that number of electrons must be added to the total valence electron count. Conversely, if the ion has a positive charge, that number must be subtracted.
The final Lewis structure for a polyatomic ion must be enclosed in brackets, with the overall charge written outside. This notation indicates that the structure represents a charged species.
To evaluate the quality of a Lewis structure, especially when multiple arrangements are possible, the formal charge of each atom is calculated. Formal charge is determined by subtracting the number of non-bonding electrons and half the number of bonding electrons from the atom’s original number of valence electrons. Structures are generally more stable if the formal charges are as close to zero as possible.
When non-zero formal charges are unavoidable, the most stable structure places the negative formal charge on the most electronegative atom. The sum of all formal charges must always equal the overall charge of the molecule or ion.
Handling Octet Exceptions and Resonance
While the octet rule is a strong guideline, certain structures require exceptions. Some molecules exhibit an incomplete octet, where the central atom is stable with fewer than eight valence electrons, as seen in Boron compounds like \(\text{BF}_3\). Boron is satisfied with only six electrons, and the incomplete octet structure is the preferred representation.
Conversely, atoms in the third period and beyond, such as Sulfur or Phosphorus, can accommodate more than eight valence electrons, known as an expanded octet. This allows the central atom to bond with five or six surrounding atoms. This exception occurs when leftover electrons are placed on the central atom, allowing it to exceed the standard octet limit.
Another complexity is resonance, which occurs when two or more valid Lewis structures can be drawn for the same molecule, differing only in the position of electrons. The actual structure is not any one of the resonance forms, but a hybrid of all of them, known as a resonance hybrid.
These resonance structures are connected by a double-headed arrow (\(\leftrightarrow\)) to indicate that the electron arrangement is delocalized across the molecule. For example, in the nitrate ion (\(\text{NO}_3^-\)), the double bond equally exists across all three Oxygen atoms.