Formal charge represents a theoretical charge assigned to an atom within a molecule, distinct from its overall ionic charge. This concept helps in understanding how electrons are distributed among atoms when they form covalent bonds, offering insights into the electron arrangement within a molecule.
What Formal Charge Means
Formal charge is a theoretical charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally between the bonded atoms. This is a conceptual tool, not an actual charge, and it helps in evaluating the distribution of electrons within a molecule. It becomes particularly useful in predicting the most stable or preferred arrangement of atoms and electrons within a molecular structure. This theoretical value aids in selecting the lowest energy structure among various possible Lewis structures for a chemical species.
Components of the Formal Charge Formula
The calculation of formal charge for an atom involves three components: the number of valence electrons the neutral atom possesses, the number of non-bonding electrons on that atom in the Lewis structure, and the number of bonding electrons shared by that atom. Valence electrons are those found in the outermost shell of a neutral atom, which can be determined from its position on the periodic table. Non-bonding electrons are the lone pair electrons located solely on a specific atom and not shared with another. Bonding electrons are those shared between two atoms within a chemical bond, with each bond contributing two electrons to this count.
Calculating Formal Charge: A Step-by-Step Guide
The formal charge for an atom in a Lewis structure uses the formula: Formal Charge = (Valence Electrons) – (Non-bonding Electrons) – (1/2 Bonding Electrons). Alternatively, it can be expressed as: Formal Charge = (Valence Electrons) – (Non-bonding Electrons) – (Number of Bonds).
To illustrate, consider the carbon dioxide (CO2) molecule. The central carbon atom forms two double bonds with two oxygen atoms. For the carbon atom, it has 4 valence electrons. In the CO2 Lewis structure with two double bonds, the carbon has no non-bonding electrons and 8 bonding electrons (4 from each double bond). Applying the formula, the formal charge for carbon is 4 – 0 – (1/2 8) = 0.
Each oxygen atom in this structure has 6 valence electrons. Each oxygen has 4 non-bonding electrons (two lone pairs) and 4 bonding electrons (from its double bond with carbon). Thus, the formal charge for each oxygen is 6 – 4 – (1/2 4) = 0.
Another example is the carbonate ion (CO3^2-). The central carbon atom is double-bonded to one oxygen and single-bonded to two other oxygens. For the carbon atom, with 4 valence electrons, it has no non-bonding electrons and 8 bonding electrons (4 from the double bond and 2 from each single bond). The formal charge for carbon is 4 – 0 – (1/2 8) = 0.
For the oxygen atom double-bonded to carbon, it has 6 valence electrons. This oxygen has 4 non-bonding electrons (two lone pairs) and 4 bonding electrons. Its formal charge calculates to 6 – 4 – (1/2 4) = 0.
Each of the two oxygen atoms single-bonded to carbon also has 6 valence electrons. These oxygens each have 6 non-bonding electrons (three lone pairs) and 2 bonding electrons. Their formal charge is 6 – 6 – (1/2 2) = -1. The sum of the formal charges (0 + 0 + (-1) + (-1)) is -2, which matches the overall charge of the carbonate ion.
Applying Formal Charge to Lewis Structures
Formal charges are used to assess the plausibility and relative stability of different Lewis structures for a molecule or ion. When multiple valid Lewis structures can be drawn, formal charges help determine which one is the most likely representation. The primary guideline is that the most stable structure is generally the one where the formal charges on all atoms are minimized, ideally resulting in zero formal charges.
If non-zero formal charges are unavoidable, the most stable structure will have negative formal charges located on the more electronegative atoms. Conversely, positive formal charges should reside on the less electronegative atoms. Additionally, structures with less separation of formal charges are considered more stable. The sum of all formal charges in a molecule must equal the overall charge of the molecule or ion, serving as a useful check for calculations.