Lewis structures serve as visual representations of molecular bonding, illustrating the arrangement of valence electrons around individual atoms. Their purpose is to help visualize the distribution of bonding and lone pair electrons within a molecule. While constructing a basic Lewis structure is a foundational skill, determining the most accurate and stable representation for a given molecule requires applying specific chemical principles. This ensures the depicted structure closely reflects the molecule’s actual electron distribution and stability.
Drawing the Initial Lewis Structure
Constructing an initial Lewis structure begins by determining the total number of valence electrons present in the molecule or ion. This involves summing the valence electrons for each atom, and for ions, adding an electron for each negative charge or subtracting one for each positive charge. Next, identify the central atom, which is typically the least electronegative element in the molecule, excluding hydrogen, as hydrogen atoms are always terminal.
Once the central atom is identified, draw single bonds connecting it to all surrounding atoms. Each single bond accounts for two shared electrons, which must be subtracted from the total valence electron count. The remaining electrons are then distributed as lone pairs, first to satisfy the octet rule for all outer atoms. If any electrons are left after the outer atoms have complete octets, they are placed on the central atom as lone pairs. If, after distributing all available electrons, the central atom still lacks a complete octet, lone pairs from adjacent outer atoms can be converted into multiple bonds (double or triple bonds).
Using Formal Charges to Refine Structures
Formal charge is a theoretical charge assigned to an atom within a molecule. This hypothetical charge helps in evaluating and selecting the most probable Lewis structure when multiple valid arrangements can be drawn for a molecule or ion. The formal charge for each atom is calculated by subtracting the number of non-bonding electrons and half the number of bonding electrons from the atom’s number of valence electrons in its neutral, isolated state.
The most favorable Lewis structure minimizes formal charges on all atoms, ideally resulting in a formal charge of zero for every atom. If non-zero formal charges are unavoidable, the preferred structure is the one with the smallest magnitudes of these charges. Furthermore, any negative formal charges should reside on the most electronegative atoms within the structure, while positive formal charges are better placed on less electronegative atoms. It is also important that the sum of all formal charges in the Lewis structure equals the overall charge of the molecule or ion. Structures with formal charges greater than two in magnitude are less stable.
Considering Octet Rule Exceptions and Electronegativity
While the octet rule is a guiding principle for many molecules, several common exceptions exist that influence the determination of the most accurate Lewis structure. Some elements, such as boron and beryllium, frequently form compounds with an incomplete octet, meaning they are stable with fewer than eight valence electrons around the central atom. For instance, boron often has only six electrons in its valence shell in compounds like BF₃. Conversely, elements from the third period and beyond can exhibit an expanded octet, accommodating more than eight electrons in their valence shell. Examples include phosphorus in PCl₅ (ten electrons) and sulfur in SF₆ (twelve electrons).
Representing Resonance Structures
In certain molecules or polyatomic ions, a single Lewis structure cannot fully describe the electron distribution, requiring resonance. Resonance occurs when multiple valid Lewis structures can be drawn for a molecule, differing only in the placement of electrons, not in the arrangement of atoms. The actual structure of such a molecule is not represented by any one of these individual forms, but rather as a hybrid or average of all contributing resonance structures.
Resonance is indicated by drawing each contributing Lewis structure separated by a double-headed arrow. Molecules that exhibit resonance often contain alternating single and multiple bonds, or a lone pair of electrons adjacent to a multiple bond, allowing for the movement of electron density. While all valid resonance structures contribute to the overall hybrid, their contributions are not always equal; structures that are more stable, often those with minimized formal charges and complete octets, contribute more significantly to the true electronic structure.