Lewis structures show how valence electrons are arranged in molecules and polyatomic ions. This cheat sheet helps students draw structures correctly, check electron counts, and compare possible bonding patterns. It is especially useful for predicting molecular stability before studying shape, polarity, and reactions.
The core ideas are counting valence electrons, placing bonds and lone pairs, applying the octet rule, and using formal charge to choose the best structure. Resonance occurs when more than one valid Lewis structure can represent the same electron arrangement. The real molecule is a resonance hybrid, not a structure switching back and forth.
Key Facts
- The total valence electron count equals the sum of valence electrons from all atoms, plus electron for each charge and minus electron for each charge.
- A single covalent bond contains shared electrons, so each single bond uses electrons from the total count.
- Most second-period atoms follow the octet rule and are most stable with valence electrons around them.
- Hydrogen follows the duet rule and is stable with valence electrons in one single bond.
- Formal charge is calculated with , where is valence electrons, is nonbonding electrons, and is bonding electrons.
- The best Lewis structure usually has the smallest formal charges and places negative formal charge on the more electronegative atom.
- Resonance structures must have the same atom positions and the same total number of valence electrons, but different electron placements.
- The bond order in a resonance hybrid can be estimated with .
Vocabulary
- Valence Electron
- A valence electron is an outer-shell electron that can participate in bonding or appear as a lone pair.
- Lewis Structure
- A Lewis structure is a diagram that shows atoms, covalent bonds, lone pairs, and sometimes formal charges.
- Lone Pair
- A lone pair is a pair of valence electrons that belongs to one atom and is not shared in a bond.
- Formal Charge
- Formal charge is the charge assigned to an atom in a Lewis structure using .
- Resonance Structure
- A resonance structure is one of two or more valid Lewis structures that differ only in electron placement.
- Resonance Hybrid
- A resonance hybrid is the real electron distribution represented by the combined effect of all valid resonance structures.
Common Mistakes to Avoid
- Forgetting to adjust for ion charge is wrong because a ion has one extra electron and a ion has one fewer electron.
- Changing atom positions between resonance structures is wrong because resonance only moves electrons, not nuclei.
- Putting more than electrons around second-period atoms like , , , or is wrong because these atoms cannot expand their octets.
- Choosing a structure with large formal charges is usually wrong because the most stable Lewis structure minimizes charge separation.
- Drawing resonance as rapid flipping is wrong because the real molecule is a resonance hybrid with delocalized electrons.
Practice Questions
- 1 Draw the Lewis structure for and state how many lone pairs are on the nitrogen atom.
- 2 Calculate the total number of valence electrons in .
- 3 For an oxygen atom with valence electrons, nonbonding electrons, and bonding electrons, calculate .
- 4 Explain why the three resonance structures of represent one resonance hybrid rather than three separate molecules.
Understanding Lewis Structures & Resonance
A practical drawing method starts by choosing a sensible skeleton. The least electronegative atom is often central because it can form several bonds. Hydrogen is never the center because it forms only one bond.
Halogens are usually outside for the same reason. Connect the atoms with single bonds first, then give outer atoms lone pairs before placing any remaining electrons on the central atom.
This order prevents a common mistake where an outer atom is left short of electrons while the center has too many. For an ion, draw square brackets around the finished structure and show its overall charge outside the brackets.
Formal charge is a bookkeeping tool, not the actual charge sitting completely on one atom. It compares an atom in a molecule with the neutral version of that atom. To find it, start with the atom's usual valence electron number.
Subtract its lone pair electrons. Then subtract half of the electrons in its bonds. A formal charge pattern helps identify weak drawings.
Large separated charges usually make a structure less favorable. A positive formal charge is usually better on a less electronegative atom, while a negative formal charge is usually better on a more electronegative atom. This explains why oxygen often carries negative formal charge more comfortably than carbon.
Resonance matters because electron density affects bond strength, bond length, and reactivity. In a set of resonance contributors, only electrons move. The atoms and their connections stay fixed.
Moving an atom creates a different structure, not another resonance form. In nitrate or carbonate ions, the double bond can be drawn in different positions. Measurements show that the equivalent bonds have equal lengths.
Their lengths fall between a typical single bond and a typical double bond. The electrons are spread across several atoms, which lowers the energy of the ion.
This spreading is called delocalization. It often makes a molecule or ion more stable than any one drawing suggests.
Students meet these ideas when studying acids, bases, organic reactions, air pollutants, and biological molecules. A conjugate base is often more stable when its negative charge can spread through resonance. That stability can make the original acid more likely to donate a hydrogen ion.
In organic chemistry, a charge next to a double bond may be delocalized across a chain of atoms. When checking work, count electrons twice, inspect every second period atom, then compare formal charges. Watch for valid exceptions.
Boron can be stable with six electrons, atoms such as phosphorus can have more than eight electrons, and radicals have an odd number of electrons. These cases do not mean the rules failed. They show that Lewis structures are useful models with limits.