Chemical Bonding and Lewis Structures

Note on scope: covalent compound naming — converting names like tetraphosphorus hexoxide to formulas P₄O₆ using Greek prefixes (mono, di, tri, tetra, …) — lives on T4: Chemical Nomenclature. This topic (T14) starts from a known molecular formula and focuses on the structure of the bond — Lewis dot pictures, octet rule, formal charge, and resonance.

Valence electrons are the outermost electrons of an atom and determine how it forms chemical bonds. A Lewis symbol represents an element by its chemical symbol surrounded by dots, one for each valence electron. For main-group elements, the number of valence electrons equals the group number (e.g., carbon in Group 14 has 4, oxygen in Group 16 has 6).

Lewis symbols illustrate ionic bond formation: sodium (one dot) transfers its electron to chlorine (seven dots), producing Na⁺ (no dots) and Cl⁻ (eight dots). For covalent bonds, atoms share electrons rather than transfer them, and Lewis structures show this sharing with lines (bonds) and dots (lone pairs).

A single shared pair is a single bond; two shared pairs form a double bond; three shared pairs form a triple bond. Each bond contributes two electrons toward satisfying the octet.

A systematic procedure produces reliable Lewis structures for molecules and polyatomic ions:

1. Count total valence electrons. Sum the valence electrons of every atom. For anions, add one electron per negative charge; for cations, subtract one per positive charge.
2. Draw a skeleton structure. Place the least electronegative atom in the center (hydrogen and fluorine are always terminal). Connect each outer atom to the center with a single bond.
3. Distribute remaining electrons as lone pairs on the terminal atoms, giving each an octet (except H, which needs only 2).
4. Place any leftover electrons on the central atom as lone pairs.
5. If the central atom lacks an octet, convert one or more lone pairs from adjacent atoms into double or triple bonds until the octet is satisfied.

For polyatomic ions, enclose the finished structure in square brackets with the overall charge as a superscript.

The octet rule states that main-group atoms tend to form bonds until they are surrounded by eight valence electrons, achieving a noble-gas electron configuration. Hydrogen is the principal exception — it needs only two electrons (a duet) to match helium.

The rule reliably predicts bonding for second-period elements (C, N, O, F). Carbon, with four valence electrons, forms four bonds (e.g., CH₄). Nitrogen forms three bonds plus one lone pair (e.g., NH₃). Oxygen forms two bonds plus two lone pairs (e.g., H₂O). Fluorine forms one bond plus three lone pairs (e.g., HF).

When single bonds alone cannot complete an octet on the central atom, atoms share additional pairs to form double bonds (as in O=C=O) or triple bonds (as in N≡N). The more electron pairs shared between two atoms, the shorter and stronger the bond.

Three categories of molecules violate the octet rule:

• Odd-electron species (free radicals) have an odd total of valence electrons, so at least one atom must have an unpaired electron. Examples include NO (11 electrons) and NO₂ (17 electrons). Free radicals are typically very reactive.
• Electron-deficient molecules have a central atom with fewer than eight electrons. Boron in BF₃ has only six electrons and beryllium in BeCl₂ has only four. These species readily accept an electron pair from a Lewis base (e.g., BF₃ + NH₃ → F₃B–NH₃).
• Expanded octets (hypervalent molecules) occur when elements in the third period or beyond use empty d orbitals to accommodate more than eight electrons. Examples: PCl₅ (10 electrons around P), SF₆ (12 around S), and XeF₂ (10 around Xe).

Second-period elements (C, N, O, F) never exceed an octet because they lack accessible d orbitals.

Formal charge (FC) is a bookkeeping tool that assigns hypothetical charges to atoms in a Lewis structure by assuming bonding electrons are shared equally:

FC = (valence electrons of free atom) − (lone-pair electrons) − ½(bonding electrons)

The sum of all formal charges must equal the overall charge of the species: zero for a neutral molecule, or the ion charge for polyatomic ions. If this sum does not match, the Lewis structure contains an error.

For example, in CO₂ drawn as O=C=O, each oxygen has FC = 6 − 4 − ½(4) = 0, and carbon has FC = 4 − 0 − ½(8) = 0. All atoms carry zero formal charge, confirming this is a good structure.

Remember that formal charge is not the actual partial charge on an atom — it is a tool for comparing alternative Lewis structures.

When more than one valid Lewis structure can be drawn for a molecule, formal charge guidelines identify the most plausible arrangement:

1. Prefer the structure in which all formal charges are zero.
2. If nonzero formal charges are unavoidable, prefer the smallest magnitudes.
3. Negative formal charges should reside on the more electronegative atoms.
4. Avoid structures with like charges on adjacent atoms.

These guidelines explain why the central atom is typically the least electronegative element. For example, CO₂ with carbon in the center and two double bonds gives all-zero formal charges, whereas placing oxygen in the center forces nonzero charges on multiple atoms. Similarly, the thiocyanate ion SCN⁻ is best drawn with carbon in the center, placing the −1 charge on the more electronegative nitrogen.

Expanded-octet resonance structures for period-3 central atoms. When the central atom can exceed an octet (P, S, Cl, Br, Xe, … — the same elements covered in “Exceptions to the Octet Rule” above), the resonance picture often includes additional structures with extra double bonds on the central atom. The phosphate ion PO₄³⁻ is the canonical example: alongside the four equivalent structures with one P=O double bond plus three P−O single bonds (FC on the doubly-bonded O = 0, on each singly-bonded O = −1, on P = 0), a structure with all four P−O bonds drawn as single bonds (FC on each O = −1, FC on P = +1) is also a contributing resonance form. Each contributing form leaves the sum of formal charges equal to the ion charge (−3), and structures with smaller magnitudes of formal charge on the central atom are typically weighted more heavily in the resonance hybrid. Both pictures together rationalize the experimental observation that all four P−O bonds in PO₄³⁻ are equivalent in length and shorter than a pure P−O single bond.

Some molecules cannot be represented by a single Lewis structure because multiple equivalent arrangements of bonds are possible. Resonance describes this situation: two or more Lewis structures (called resonance forms) are drawn with a double-headed arrow between them, and the true electron distribution is their average, called the resonance hybrid.

The nitrite ion NO₂⁻ can be drawn with the double bond on either oxygen. Experiments show both N–O bonds are identical in length and strength — intermediate between a single and double bond. The molecule does not “flip” between forms; the hybrid is the only real structure.

The carbonate ion CO₃²⁻ has three equivalent resonance forms. Each C–O bond has a bond order of 1⅓ (one-third double-bond character), and all three bonds are experimentally identical. Resonance is common whenever a molecule has multiple equivalent positions for a double bond, and it generally stabilizes the molecule by delocalizing electron density.

Follow this decision strategy to avoid the most common Lewis-structure errors:

1. Count total valence electrons first. For polyatomic ions, add electrons for negative charge or subtract for positive charge.
2. Place the least electronegative atom in the center (H is always terminal).
3. Distribute remaining electrons as lone pairs to satisfy octets, starting with terminal atoms.
4. If the central atom lacks an octet, convert lone pairs on adjacent atoms into multiple bonds.
5. Calculate formal charges on each atom. The best structure minimizes formal charges and places any negative formal charge on the more electronegative atom.

Common mistakes: miscounting total valence electrons (especially for ions), placing the most electronegative atom in the center, forgetting to check for expanded octets on period-3+ elements, drawing multiple bonds when single bonds already satisfy all octets, and confusing formal charge with oxidation state.