8.1 Lewis Structures, Resonance & Formal Charge

Drawing Lewis Structures

A Lewis structure (Lewis electron dot formula) shows the arrangement of valence electrons in a molecule or ion as bonding pairs shared between atoms and lone pairs localized on a single atom. The MCAT expects you to draw a correct structure quickly and then use it to predict geometry, polarity, and reactivity.

A reliable method:

1. Count total valence electrons — sum the group number of each neutral atom, add one electron per unit of negative charge, and subtract one electron per unit of positive charge.
2. Choose the central atom, usually the least electronegative atom present (hydrogen is never central; it always forms exactly one bond).
3. Connect every outer atom to the central atom with a single bond, then distribute the remaining electrons as lone pairs to satisfy octets, starting with the terminal (outer) atoms.
4. If the central atom still lacks an octet after this, convert a lone pair on a terminal atom into a second or third shared pair, forming a double or triple bond, until every atom (where possible) has eight electrons.

Exceptions to the octet rule are a favorite MCAT trap:

• Electron-deficient atoms: boron and beryllium compounds, such as boron trifluoride (BF3) and beryllium chloride (BeCl2), are stable with fewer than eight electrons around the central atom because these elements have only three or two valence electrons to contribute.
• Expanded octets: elements in period 3 and beyond — phosphorus, sulfur, chlorine, and others — can exceed eight electrons around a central atom, as in phosphorus pentachloride (PCl5, 10 electrons) and sulfur hexafluoride (SF6, 12 electrons). Second-row elements (carbon, nitrogen, oxygen, fluorine) can never expand their octet because they have no accessible d-orbitals.
• Odd-electron species: radicals such as nitric oxide (NO) and nitrogen dioxide (NO2) have an odd total number of valence electrons, so at least one atom cannot have a complete octet.

Formal Charge

Formal charge (FC) is an electron-counting device used to judge which of several valid Lewis structures is the most realistic representation of a molecule. It is calculated atom by atom as:

FC = (valence electrons of the free atom) − (nonbonding electrons) − 1/2(bonding electrons)

Worked example — ozone (O3): Ozone has 18 valence electrons (6 + 6 + 6). One valid structure places a single bond between the central oxygen and one terminal oxygen, and a double bond between the central oxygen and the other terminal oxygen. Calculating formal charge on each atom:

• Singly bonded terminal O: 6 − 6 (three lone pairs) − 1/2(2 bonding electrons) = 6 − 6 − 1 = −1
• Central O: 6 − 2 (one lone pair) − 1/2(6 bonding electrons) = 6 − 2 − 3 = +1
• Doubly bonded terminal O: 6 − 4 (two lone pairs) − 1/2(4 bonding electrons) = 6 − 4 − 2 = 0

These formal charges (−1, +1, 0) sum to zero, matching ozone's overall neutral charge — a mandatory check for any structure. Because ozone's two oxygen-oxygen bonds are experimentally identical in length, this single structure is incomplete on its own; it must be combined with its mirror-image resonance structure (double bond on the opposite side) to represent the true, delocalized molecule, covered next.

Rules for choosing the best Lewis structure:

1. Minimize the number of formal charges and keep them as close to zero as possible.
2. When charge separation cannot be avoided, place any negative formal charge on the more electronegative atom.
3. Avoid placing like formal charges (two positives or two negatives) on adjacent atoms, since this is electrostatically unfavorable.

Formal charge is a bookkeeping approximation that assumes perfectly equal sharing of bonding electrons — it is not the same as the real, unequal distribution of electron density that arises from electronegativity differences (partial charge), discussed later in this section.

Resonance Structures

Some molecules and ions cannot be accurately represented by a single Lewis structure. When two or more valid Lewis structures exist for the same arrangement of atoms — differing only in the placement of electrons, never in the positions of the nuclei — they are called resonance structures, conventionally connected by a double-headed arrow. The actual molecule is a single resonance hybrid, an average of the contributing structures, not a mixture that flips back and forth between them.

Classic examples:

• Ozone (O3): two resonance structures place the double bond on alternating sides; the true molecule has two equivalent oxygen-oxygen bonds, each with a bond order of 1.5 (partway between a single and a double bond).
• Carbonate ion (CO3^2-): three resonance structures distribute one double bond among three equivalent carbon-oxygen positions; each real bond has a bond order of 4/3 and a length intermediate between a typical single bond (about 143 pm) and double bond (about 123 pm).
• Benzene (C6H6): six pi electrons delocalize around the entire six-membered ring rather than forming three isolated double bonds, giving all six carbon-carbon bonds identical length (about 139 pm) and roughly 36 kilocalories per mole of extra stability (resonance/delocalization energy) compared to a hypothetical molecule with three isolated double bonds.

Curved-arrow notation tracks how one resonance structure converts into another on paper: an arrow's tail starts at an electron-rich site (a lone pair or a pi bond) and its head points to where that electron pair moves, without any atoms themselves moving.

Resonance versus tautomers — a frequent trap: Resonance structures have identical atom connectivity and differ only in electron placement; no bonds to atoms (especially hydrogens) are broken or formed. Tautomers, such as the keto and enol forms of a carbonyl compound, are different constitutional isomers — a hydrogen atom and a double bond genuinely relocate — that interconvert through an actual chemical reaction and exist in true equilibrium. Resonance structures are not isomers of each other and cannot be isolated individually; only one real molecule exists.

Biological relevance: The peptide (amide) bond linking amino acids is stabilized by resonance between a lone pair on the amide nitrogen and the adjacent carbonyl pi bond, giving the carbon-nitrogen bond partial double-bond character. This resonance restricts rotation around the peptide bond and forces it to be planar, a geometric constraint fundamental to how alpha helices and beta sheets fold. Carboxylate resonance in deprotonated fatty acids and in the aspartate and glutamate side chains similarly delocalizes negative charge over both oxygens, stabilizing the anion and lowering its pKa relative to a simple alcohol.

Lewis Acids and Bases

The Bronsted-Lowry definition classifies acids and bases by proton transfer, but the Lewis definition is broader and classifies them by electron-pair transfer: a Lewis acid accepts an electron pair (it has an empty or accessible low-energy orbital), and a Lewis base donates an electron pair (it has an available lone pair or pi bond).

Every Bronsted-Lowry acid-base pair is technically also a Lewis pair — a proton (H+) is the simplest possible Lewis acid, since it has a completely empty 1s orbital — but the Lewis framework also captures many important reactions that involve no protons at all:

• Boron trifluoride and ammonia: BF3 + NH3 → F3B-NH3. Boron has only six valence electrons around it in BF3 and an empty p orbital, making it electron-deficient and a strong Lewis acid; ammonia's nitrogen lone pair makes it the Lewis base.
• Metal cations in biology: Fe3+/Fe2+ in the heme group of hemoglobin and myoglobin, Zn2+ in the active site of carbonic anhydrase and many other metalloenzymes, and Mg2+ coordinating the phosphate groups of ATP all act as Lewis acids, accepting electron density donated by ligand lone pairs (histidine nitrogen, carboxylate oxygen, water oxygen).
• Carbonyl chemistry: the carbon of a ketone or aldehyde carbonyl group is electrophilic and behaves as a Lewis acid toward nucleophilic Lewis bases such as cyanide or amines, a pattern that underlies countless organic reaction mechanisms tested on the MCAT.

Common trap: do not assume 'acid' always means 'proton donor.' A compound with no acidic hydrogens at all, such as BF3 or the carbonyl carbon of acetone, can still function as a Lewis acid by accepting an electron pair.

Partial Ionic Character and Dipole Moment

No covalent bond between two different elements involves perfectly equal electron sharing — the more electronegative atom pulls shared electron density toward itself, creating partial ionic character. The size of this effect tracks the electronegativity difference (ΔEN) between the bonded atoms on the Pauling scale, which rises across a period and falls down a group. Fluorine is the most electronegative element (3.98), followed by oxygen (3.44), chlorine (3.16), and nitrogen (3.04); carbon (2.55) and hydrogen (2.20) sit in the middle of the scale.

• ΔEN near 0: essentially nonpolar covalent (identical atoms bonded together, such as O-O or N-N, are always exactly 0; a C-H bond, ΔEN ≈ 0.35, is treated as effectively nonpolar).
• ΔEN roughly 0.5 to 1.7: polar covalent — this range includes most biologically important bonds, including O-H, N-H, C-O, C=O, and C-N.
• ΔEN above roughly 1.7: the bond is generally considered predominantly ionic.

Each polar bond carries a bond dipole moment, a vector pointing from the partially positive atom (δ+) toward the partially negative atom (δ−), with magnitude μ = q × r (the partial charge times the distance separating the charges), measured in debyes (D).

A molecule's overall dipole moment is the vector sum of every individual bond dipole — which is why molecular geometry, covered in the next section, is essential for predicting polarity, and why the MCAT loves this trap: a molecule can contain highly polar bonds yet be completely nonpolar overall if those bond dipoles are arranged symmetrically enough to cancel.

Notice that replacing one chlorine of CCl4 with a hydrogen to form chloroform (CHCl3) breaks the molecule's symmetry, so the bond dipoles no longer cancel and the molecule becomes polar overall, even though it looks superficially similar to carbon tetrachloride.

Finally, remember that formal charge and partial charge are different tools measuring different things. Formal charge is a bookkeeping calculation that assumes perfectly equal electron sharing between bonded atoms, useful for comparing candidate Lewis structures. Partial charge (δ+/δ−) reflects the real, unequal distribution of electron density predicted by electronegativity, and it is partial charge — not formal charge — that determines a molecule's actual polarity and dipole moment.