Lewis structures are visual representations of the bonding within a molecule or a polyatomic ion, illustrating the arrangement of atoms and electrons. Developed by Gilbert N. Lewis in 1916, these diagrams primarily focus on valence electrons, which are the electrons available for chemical bonding. Lewis structures provide a simple, two-dimensional understanding of how atoms share or transfer these outer-shell electrons to achieve a stable electronic configuration. By showing the connectivity between atoms and the placement of non-bonding electrons, they help predict aspects like the number of bonds an atom will form and the overall molecular shape.
The Fundamental Rules of Construction
The foundation of drawing Lewis structures rests on understanding valence electrons. The number of valence electrons an atom possesses determines its bonding behavior and is typically found by looking at the atom’s group number on the periodic table. Atoms tend to share electrons to attain the most stable arrangement, usually involving a full outer shell similar to noble gases.
This drive for stability is formalized by the Octet Rule, which states that most atoms strive to be surrounded by eight valence electrons, either through sharing in bonds or as non-bonding lone pairs. The Duet Rule is an exception applying to hydrogen and helium, as they only require two electrons to complete their outer shell.
In a Lewis structure, a single covalent bond is represented by a line between two atoms, signifying one shared pair (two electrons). A double bond involves two lines (four shared electrons), and a triple bond uses three lines (six shared electrons). Electrons that do not participate in bonding are known as lone pairs, shown as pairs of dots placed next to the atomic symbol.
A Step-by-Step Guide to Drawing
The process of constructing a Lewis structure begins by calculating the total number of valence electrons contributed by all atoms. For example, in ammonia (\(\text{NH}_3\)), nitrogen contributes five valence electrons and the three hydrogen atoms contribute one each, totaling eight electrons that must be accounted for.
Next, identify the central atom, which is typically the least electronegative atom; hydrogen is always a terminal atom. Draw a single bond between the central atom and each surrounding atom, using two valence electrons per bond. In \(\text{NH}_3\), this uses six electrons, leaving two remaining.
The remaining electrons are distributed as lone pairs, first to the terminal atoms until their octets (or duets for hydrogen) are satisfied. Since hydrogen atoms satisfy the Duet Rule with their single bonds, the remaining two electrons are placed on the central nitrogen atom.
The final step requires checking that all atoms, especially the central one, have satisfied the Octet Rule. If the central atom is deficient, a lone pair from a terminal atom must be converted into an additional bond, forming a double or triple bond to ensure the central atom reaches its octet.
Handling Complex Molecules and Exceptions
When drawing structures for charged species, known as polyatomic ions, the initial electron count must be adjusted to account for the overall charge. For a negatively charged ion, one electron is added for every unit of negative charge, while a positive charge requires subtracting one electron. The final Lewis structure for any polyatomic ion must be enclosed in brackets with the charge written outside.
Some molecules naturally deviate from the Octet Rule, exhibiting incomplete octets where the central atom is stable with fewer than eight valence electrons. Elements from Group 13, such as boron and aluminum, often form stable compounds where the central atom has only six electrons. This electron deficiency signifies a tendency to react with molecules that can donate a pair of electrons.
Conversely, elements in the third period and beyond, like sulfur or phosphorus, can accommodate more than eight electrons, a phenomenon called an expanded octet. This is possible because these larger atoms have access to empty \(d\) orbitals, allowing them to form more than four bonds. Expanded octets are accepted if they minimize the formal charge on the atoms.
In cases where a single Lewis structure cannot accurately describe the bonding, the concept of resonance structures is employed. Resonance occurs when multiple valid Lewis structures can be drawn for the same molecule, differing only in the placement of electrons, not the atoms. The actual structure is considered a hybrid, or average, of all possible structures.