Difficulty: Introductory / Foundational
Prerequisites: Part 1 of these notes (atomic structure, electron configuration, orbital shapes). You need to be comfortable with valence electrons and the octet rule before tackling Lewis structures and formal charge.
Once you know how electrons are arranged in atoms, the next question is: how do atoms use those electrons to bond? This section covers the two main bonding types (ionic and covalent), how to draw Lewis structures that show bonding and lone pairs, how to assign formal charges, and how resonance structures describe molecules that a single Lewis structure cannot capture. These skills are the bread and butter of organic chemistry; you will use them in every chapter that follows.
Ionic bonds form when atoms transfer electrons; covalent bonds form when atoms share them. Lewis structures map out which atoms are bonded and where the lone pairs sit. Formal charge tells you where charge accumulates when atoms share electrons unevenly. When one Lewis structure is not enough to describe a molecule, you draw resonance structures; the real molecule is a weighted average (hybrid) of all of them and is more stable than any single contributor.
Ionic bond
A bond formed by the electrostatic attraction between a positively charged ion (cation) and a negatively charged ion (anion). Ions form when atoms transfer electrons to achieve a noble-gas electron configuration. Think of it as: one atom gives, the other takes, and opposite charges hold them together.
Covalent bond
A bond formed when two atoms share one or more pairs of electrons. In organic chemistry, nearly all bonds are covalent. Think of it as: atoms hold hands by sharing electrons rather than giving them away.
Lewis structure (Lewis dot structure)
A diagram showing all valence electrons in a molecule as bonding pairs (lines) and lone pairs (dots). The goal is to give every atom a complete octet (or duet for hydrogen).
Bonding electrons
Electrons shared between two atoms in a covalent bond. Each single bond contains two bonding electrons.
Nonbonding electrons (lone pairs)
Valence electrons on an atom that are not shared with another atom. They sit as pairs in orbitals and affect molecular geometry and reactivity.
Formal charge
The charge assigned to an atom in a Lewis structure, calculated as: (valence electrons of the free atom) minus (nonbonding electrons) minus (half the bonding electrons). It is a bookkeeping device, not always the actual charge on the atom.
Resonance structures
Two or more valid Lewis structures for the same molecule that differ only in the placement of electrons (not atoms). No single structure fully describes the molecule.
Resonance hybrid
The actual structure of a molecule described by resonance, which is a weighted average of all contributing resonance structures. The hybrid is always more stable than any individual contributor. Think of it as: the real molecule is a blend, not a flip between the structures.
Bond order
The number of bonding electron pairs shared between two atoms. A single bond has order 1, a double bond has order 2, a triple bond has order 3. In resonance hybrids, bond orders can be fractional (e.g. 1.5 for benzene).
Ionic bonds form between atoms with very different electronegativities (typically a metal and a nonmetal). One atom loses electrons, the other gains them, creating cations and anions held together by electrostatic attraction.
The strength of an ionic bond depends on the product of the charges and the distance between the ions (Coulomb's law). CaCl₂ has a higher lattice energy (2195 kJ/mol) than KCl (715 kJ/mol) because Ca²⁺ has a larger charge and a smaller ionic radius than K⁺.
Covalent bonds form between atoms of similar electronegativity (typically two nonmetals). Atoms share electrons rather than transferring them.
Both ionic and covalent bonding follow the octet rule: each atom aims for eight valence electrons (two for hydrogen).
Count total valence electrons for the molecule or ion. For anions, add one electron per negative charge. For cations, subtract one electron per positive charge.
Place bonding pairs between atoms (each bond = 2 electrons), then distribute remaining electrons as lone pairs to satisfy octets.
Hydrogen always makes one bond. Carbon makes four bonds. Nitrogen makes three bonds and has one lone pair. Oxygen makes two bonds and has two lone pairs. Halogens make one bond and have three lone pairs.
Example: chloroform (CHCl₃) has carbon at the centre making four single bonds, one to H and three to Cl, each Cl carrying three lone pairs.
Formal charge = (valence electrons of free atom) − (lone pair electrons) − (½ × bonding electrons).
In a correct Lewis structure, the sum of all formal charges equals the overall charge of the molecule or ion.
Formal charge is a bookkeeping tool. The actual charge distribution depends on electronegativity. For example, in the ammonium ion (NH₄⁺), the formal charge of +1 sits on nitrogen, but because nitrogen is more electronegative than hydrogen, most of the actual positive charge resides on the hydrogens.
When a molecule cannot be described by a single Lewis structure, you draw two or more resonance structures connected by a double-headed arrow (↔).
Resonance structures differ only in electron placement; the atoms stay in the same positions.
The real molecule (the resonance hybrid) is a weighted average of all contributors and is more stable than any single one.
Example: the allyl anion [CH₂=CH-CH₂]⁻ has two equivalent resonance structures. Each terminal carbon bears −0.5 charge, and each C-C bond has a bond order of 1.5.
Example: benzene has two equivalent resonance structures in which single and double bonds alternate. Each C-C bond in the hybrid has a bond order of 1.5.
The carbonate ion (CO₃²⁻) has three equivalent resonance structures. Each C-O bond order is 4/3, and each oxygen carries −2/3 of a unit charge.
Students often think resonance structures represent the molecule flipping back and forth between forms. They do not. The molecule has one fixed structure (the hybrid) at all times; the individual resonance structures are just our imperfect way of drawing it.
Formal charge is frequently confused with actual charge. A nitrogen atom with a formal charge of +1 does not necessarily carry a full +1 charge in reality; electronegativity differences redistribute the actual electron density.
Students sometimes forget to count lone pairs when checking the octet rule. An atom's octet includes both bonding and nonbonding electrons. A nitrogen with three bonds and one lone pair has a full octet (3×2 + 2 = 8).
Some students assume that all atoms must obey the octet rule. Boron commonly has only six electrons (as in BF₃), and third-row elements like sulfur and phosphorus can exceed eight.
⚠️ Drawing correct Lewis structures (including lone pairs and formal charges) is tested on virtually every organic chemistry exam. Speed and accuracy here save time on harder problems.
⚠️ Formal charge calculations will appear repeatedly. Know the formula cold: valence electrons − lone pair electrons − ½(bonding electrons).
⚠️ Resonance is central to understanding stability, acidity, basicity, and reaction mechanisms throughout the course. You will be asked to draw all significant resonance structures and predict the hybrid.
⚠️ Bond order in resonance hybrids (e.g. benzene's 1.5, carbonate's 4/3) is a common calculation question.
True or false: In a covalent bond, electrons are transferred from one atom to another.
False. Covalent bonds involve sharing electrons; ionic bonds involve transfer.
Fill in the blank: The formal charge on nitrogen in the ammonium ion (NH₄⁺) is ______.
+1.
True or false: Resonance structures represent the molecule rapidly interconverting between different forms.
False. The molecule is always the hybrid; individual resonance structures are a drawing convenience.
Fill in the blank: Each C-C bond in benzene has a bond order of ______.
1.5.
True or false: Boron in BF₃ violates the octet rule because it has only six electrons around it.
True. Boron is a well-known exception to the octet rule.
Q: Draw the Lewis structure of diazomethane (CH₂N₂) and assign formal charges to each atom.
A: Two valid structures exist. In one, H₂C=N=N with a double bond from C to the first N and a double bond from the first N to the second N; the first N has a formal charge of +1 and the terminal N has a formal charge of −1, carbon is neutral. In the other, H₂C-N≡N with a single bond from C to N and a triple bond between the two nitrogens; carbon carries −1, the central N carries +1, and the terminal N is neutral.
Q: The carbonate ion (CO₃²⁻) has three equivalent resonance structures. What is the bond order of each C-O bond, and what is the charge on each oxygen in the hybrid?
A: Each C-O bond is a double bond in one structure and a single bond in the other two. Bond order = (2 + 1 + 1) / 3 = 4/3. Each oxygen carries −2/3 of a unit charge.
Q: Explain why the ionic bond in CaCl₂ is stronger than the ionic bond in KCl.
A: The interaction energy is proportional to the product of the charges on the ions. Ca²⁺ has twice the charge of K⁺, so the electrostatic attraction to Cl⁻ is roughly doubled. Additionally, Ca²⁺ has a smaller ionic radius than K⁺, which shortens the distance between the charges and further strengthens the interaction.
Q: In the Lewis structure of the methyl anion (CH₃⁻), how many bonds does carbon make, and does it satisfy the octet rule?
A: Carbon makes three bonds to hydrogen (6 bonding electrons) and has one lone pair (2 electrons), for a total of 8 electrons. It has a complete octet and carries a formal charge of −1.
Q: Why does formal charge on oxygen in formate (HCO₂⁻) not accurately reflect the actual charge distribution?
A: In one resonance structure, one oxygen has a formal charge of −1 and the other has 0, but in the other resonance structure these roles are reversed. The hybrid distributes −0.5 on each oxygen. Even this is approximate because oxygen is more electronegative than carbon, so oxygen pulls additional electron density beyond what formal charge suggests.
Resonance is the gateway to understanding conjugation, aromaticity (Chapter 15+), and the stability of reaction intermediates like carbocations and carbanions. Formal charge directly feeds into predicting nucleophiles and electrophiles in reaction mechanisms. The distinction between ionic and covalent character becomes important when you study leaving groups, solvents, and salt formation in later chapters.
ionic bond, covalent bond, Lewis structure, Lewis dot structure, bonding electrons, nonbonding electrons, lone pair, formal charge, resonance, resonance structures, resonance hybrid, bond order, octet rule, expanded octet, Coulomb's law, lattice energy, carbonate ion, allyl anion, benzene resonance, diazomethane, ammonium ion, methyl anion, electrostatic attraction