Lewis Structures

When two atoms of similar electronegativity come together, transferring an electron to form ions costs more energy than it would release — so instead of transferring electrons, the atoms share them. This is covalent bonding, and Lewis structures are the standard way to track exactly how those shared and unshared electrons are arranged. This subtopic covers bond order, the three descriptors of a covalent bond, coordinate covalent bonds, the step-by-step rules for drawing a Lewis structure, formal charge, and resonance.

Key Takeaways

• Covalent bonds share electron pairs; bond order counts the shared pairs (1 for single, 2 for double, 3 for triple).

• Bond order determines bond length (shorter with higher order) and bond energy (higher with higher order).

• Polarity arises from an electronegativity difference between bonded atoms; the dipole moment (p = qd, measured in Debye) quantifies the resulting charge separation.

• A coordinate covalent bond forms when one atom donates both bonding electrons — the basis of Lewis acid–base chemistry (e.g., NH₃ + H⁺ → NH₄⁺).

• Drawing a Lewis structure: sum valence electrons, determine connectivity (least electronegative atom central), distribute electrons (bonds, then outer octets, then central atom), check the central atom's octet, and verify the total electron count.

• Formal charge = (valence electrons) − (lone pair electrons) − ½(bonding electrons); it helps evaluate how reasonable a given Lewis structure is.

• Resonance occurs when multiple valid Lewis structures describe the same molecule; the real structure is a hybrid. Rank contributing structures by: full octets first, then minimized formal charges, then negative charge on the more electronegative atom, then equal contribution if still tied.

Covalent Bonds: Single, Double, Triple, and Bond Order

When an atom needs to fill its valence shell, one bond isn't always enough — it can form multiple bonds with the same or different atoms. Sharing one pair of electrons is a single bond; sharing two pairs is a double bond; sharing three pairs is a triple bond. The more electron pairs two atoms share, the stronger the bond between them.

Bond order is the number of electron pairs shared between two atoms: a single bond has a bond order of 1, a double bond a bond order of 2, and a triple bond a bond order of 3.

Describing a Covalent Bond: Length, Energy, and Polarity

Three properties describe any covalent bond: bond length, bond energy, and polarity.

Bond length is the average distance between the two bonded nuclei. As the number of shared electron pairs increases, the two atoms are pulled closer together — so for a given pair of atoms, a triple bond is shorter than a double bond, which is shorter than a single bond.

Bond energy is the energy required to break a bond, separating the atoms into isolated, gaseous states. More shared electron pairs means a stronger pull between the nuclei, so triple bonds have the greatest bond energy and single bonds the least.

Polarity describes how evenly the electron pair is shared. When the two bonded atoms have different electronegativities, the atom with the higher electronegativity pulls more of the electron density toward itself. This creates a dipole, with the positive end at the less electronegative atom and the negative end at the more electronegative atom. In HCl, for example, chlorine (more electronegative) carries a partial negative charge, while hydrogen carries a partial positive charge — this is a polar bond, and HCl is a polar molecule.

The dipole moment quantifies this separation of charge. It's a vector quantity given by:

p = qd

where p is the dipole moment, q is the magnitude of the partial charge, and d is the displacement vector separating the two partial charges. The dipole moment vector points from the positive charge to the negative charge, and dipole moments are measured in Debye units.

Coordinate Covalent Bonds

A coordinate covalent bond is a covalent bond in which one atom contributes both bonding electrons — a lone pair from one atom is donated to another atom that needs electrons to complete a bond. Once the bond forms, it's indistinguishable from any other covalent bond; the distinction only matters for keeping track of valence electrons and formal charges.

Coordinate covalent bonds are central to Lewis acid–base chemistry: a Lewis base (electron pair donor) forms a coordinate covalent bond with a Lewis acid (electron pair acceptor). Classic examples include ammonia accepting a proton to form ammonium (NH₃ + H⁺ → NH₄⁺) and water accepting a proton to form hydronium (H₂O + H⁺ → H₃O⁺) — in both cases, the nitrogen or oxygen lone pair supplies both electrons in the new bond.

Drawing Lewis Structures

Covalent bonding involves sharing electron pairs between atoms. The electrons involved in a bond are bonding electrons; those not involved are nonbonding electrons, or lone pairs. A Lewis structure (or Lewis dot diagram) tracks both: it shows an element's chemical symbol, with bonding electrons drawn as lines between atoms and lone pairs drawn as pairs of dots on individual atoms. The number of valence electrons an atom starts with matches its group number on the periodic table (excluding transition metals).

To draw a Lewis structure:

1. Sum the valence electrons of all atoms in the molecule or ion. For a polyatomic ion, add one electron per negative charge and subtract one electron per positive charge.

2. Determine connectivity. The central atom is usually listed first, and is usually the least electronegative atom (hydrogen is never central — it can only form one bond). In oxyacids like HNO₃ or H₃PO₄, each hydrogen attaches to an oxygen, and the oxygens attach to the central atom. Some molecular formulas directly show bonding order, such as HCN or HOCl.

3. Distribute the electrons, in this order: first connect the atoms with single bonds (each bond uses 2 electrons); then complete the octets of the outer atoms; then place any leftover electrons on the central atom, in pairs.

4. Check the central atom's octet. If it has fewer than eight electrons, convert one or more of the outer atoms' lone pairs into double or triple bonds with the central atom until it reaches an octet.

5. Final check. Confirm the structure uses the exact total number of electrons calculated in step 1 — an incorrect electron count is the most common Lewis structure error.

Worked example — HCN. Sum the valence electrons: H contributes 1, C contributes 4, N contributes 5, for a total of 10. Hydrogen must occupy a terminal position, and between carbon and nitrogen, carbon is the less electronegative atom, so carbon takes the central position: H–C–N. Drawing single bonds first uses 4 electrons and gives hydrogen its full 2-electron duet. Distributing the remaining 6 electrons as lone pairs on nitrogen would leave carbon with only 4 electrons around it — short of an octet. So two of nitrogen's lone pairs shift into bonding pairs, forming a triple bond between carbon and nitrogen. The final structure, H–C≡N, gives carbon and nitrogen full octets and hydrogen its duet, using all 10 electrons.

Formal Charge

A Lewis structure can assign an atom a different number of valence electrons than it has as a neutral, isolated atom, because of how bonding and lone pairs are distributed. This difference is the atom's formal charge, calculated as:

Formal charge = (valence electrons in the free atom) − (lone pair electrons) − ½(bonding electrons)

Worked example — formal charge on nitrogen in NH₄⁺. Nitrogen normally has 5 valence electrons. In the NH₄⁺ Lewis structure, nitrogen forms four single bonds (to four hydrogens) and has zero lone pairs. That's 8 bonding electrons and 0 lone-pair electrons around nitrogen. Formal charge = 5 − 0 − ½(8) = 5 − 4 = +1. This matches the overall +1 charge on the ammonium ion, since nitrogen is the only atom in NH₄⁺ with a nonzero formal charge.

Resonance Structures

Lewis structures are simple models, and they don't always capture a molecule's bonding perfectly. Resonance occurs when more than one valid Lewis structure can be drawn for the same molecule — same arrangement of atoms, different placement of electrons. Resonance structures are connected by a double-headed arrow, and the molecule's actual electron distribution is a resonance hybrid: a composite of all the valid resonance structures, not any single one of them.

Worked example — [NO₃]⁻. Nitrogen sits at the center, bonded to three oxygen atoms. The ion has 24 valence electrons total: 5 from nitrogen, 6 from each of the three oxygens (18), plus 1 for the ion's negative charge. Three equivalent resonance structures can be drawn, each with one N=O double bond and two N–O single bonds — the only difference between them is which oxygen holds the double bond. Because all three contribute equally, the actual structure is a hybrid where the electrons are delocalized across all three oxygens: every N–O bond is identical, with a bond order between that of a single bond and a double bond.

Resonance structures don't always contribute equally, though. When they don't, use these rules — in order of importance — to identify the most significant contributor:

1. Satisfying the octet rule. Structures where more atoms have complete octets are more stable and contribute more to the hybrid.

2. Minimizing formal charges. Structures with small or no formal charges are preferred over structures with large formal charges.

3. Placing charges on the right atoms. A negative formal charge is more stable on a more electronegative atom (like oxygen); a positive formal charge is more stable on a less electronegative atom.

4. Equal contribution. If rules 1–3 don't distinguish between the structures — as with the three NO₃⁻ resonance forms above — they all contribute equally to the hybrid.

Common MCAT Mistakes

• Skipping the final electron-count check. The most common Lewis structure error is using the wrong total number of electrons — always re-verify the structure uses exactly the number calculated in step 1, no more, no less.

• Assuming formal charge equals real, measurable charge. Formal charge is a bookkeeping tool that assumes perfectly equal sharing of every bonding pair — it's a way to evaluate a structure's reasonableness, not a direct measurement of actual electron density.

• Placing hydrogen as a central atom. Hydrogen can only form one bond (it needs just 2 electrons for a full duet), so it can never sit in a central, multiply-bonded position in a Lewis structure.

• Treating unequal resonance contributors as equal. Resonance structures are only equal contributors when they tie on all three ranking criteria (octets, formal charge magnitude, charge placement); otherwise, one structure dominates the hybrid more than the others.

MCAT-Style Concept Check

Question: In the Lewis structure of CO₂ (O=C=O), what is the formal charge on the central carbon atom?

• A) −1

• B) 0

• C) +1

• D) +2

Answer: B

Explanation: Carbon has 4 valence electrons as a free atom. In CO₂, carbon forms two double bonds (to each oxygen) and has zero lone pairs, giving it 8 bonding electrons and 0 lone-pair electrons. Formal charge = 4 − 0 − ½(8) = 4 − 4 = 0. A formal charge of zero on the central atom is one sign of a well-drawn, low-energy Lewis structure.

FAQ

What is bond order, and how does it relate to bond length and energy?

Bond order is the number of electron pairs shared between two bonded atoms (1 for single, 2 for double, 3 for triple). Higher bond order means the atoms are pulled closer together (shorter bond length) and held more tightly (higher bond energy).

What's the difference between a regular covalent bond and a coordinate covalent bond?

Both are covalent bonds where an electron pair is shared. In a regular covalent bond, each atom contributes one electron to the shared pair; in a coordinate covalent bond, one atom donates both electrons of the pair, as when ammonia's nitrogen lone pair forms the fourth N–H bond in NH₄⁺.

How do you calculate formal charge?

Formal charge = (valence electrons in the free atom) − (lone pair electrons) − ½(bonding electrons). It's used to compare candidate Lewis structures — the structure(s) with the smallest formal charges, correctly placed, are generally the most accurate representation.

What makes resonance structures equal contributors to the hybrid?

They're equal contributors only when they tie on all three ranking criteria: all satisfy the octet rule equally well, all have equally minimized formal charges, and any charges are placed on equally appropriate atoms. The three NO₃⁻ resonance structures meet all three conditions, so they contribute equally.