Lewis Structures, Resonance & Bond Models · Grade 11 Science
Chapter: Bonding, Formulas & Molecular StructureLewis Structures, Resonance & Bond Models
Learning goals
- I can count valence electrons and check a Lewis structure’s total charge.
- I can distinguish resonance, electron-domain geometry, and molecular shape.
Readings for this unit
Before this unit
Valence electrons, ionic charge, and covalent bonding.
Use the detailed parent plan: two core 50–60-minute sessions per focused unit, with a third assigned for selected calculation/modeling topics. First study the core explanation and Parts 1–2 with their figures/checks; next study Parts 3–4, the whole-unit application, worked cases, varied practice and exit check. Attempt reasoning before opening guidance. Extend time when actual understanding requires it; investigations and chapter assessments have separate sessions.
Start with something to explain
Do two resonance drawings show a molecule changing back and forth?
Illustrative evidence case. Two ozone contributors place the formal double bond on opposite sides without moving atoms.
Notice and wonder
Record two observations using the stated quantities or relationships. Keep observations separate from possible causes. Write one question that a new measurement could help answer.
Do two resonance drawings show a molecule changing back and forth?
Your first model
Keep atom positions fixed and track only electron assignments.
Build the language from the idea
- lone pair
- A valence electron pair not represented as a bond.
- formal charge
- An electron-counting assignment, distinct from partial charge.
- resonance
- Multiple electron drawings representing delocalized bonding in one structure.
A Lewis structure is a model of valence electrons: dots represent lone electrons and lines represent shared pairs. Count all valence electrons first; add electrons for negative charge and subtract for positive charge. Identify a central atom, usually a less electronegative non-hydrogen atom, and connect terminal atoms with single bonds.
Distribute remaining electrons to complete terminal octets, with hydrogen needing only two. Put leftovers on the central atom; if its octet is incomplete, consider a double or triple bond. Recount electrons. This procedure explains CO₂ with two double bonds and H₂O with two O–H bonds and two lone pairs on oxygen.
The octet rule is a useful pattern, not a universal law. Hydrogen follows a duet; boron compounds can have incomplete octets; odd-electron molecules exist; some third-period and heavier atom structures are conventionally drawn with expanded valence shells. A model should be applied within its scope.
Formal charge is bookkeeping: valence electrons minus nonbonding electrons minus half of bonding electrons. It is not the same as measured partial charge. For NH₄⁺, nitrogen has 5 − 0 − 4 = +1 and each hydrogen has 0 formal charge; the total is +1.
Resonance structures use the same atom arrangement but different electron placement. The actual bonding is described by a delocalized structure, not a molecule rapidly switching between drawings. In nitrate, three equivalent N–O arrangements contribute; the real bonds are equivalent in the idealized ion.
Use Lewis structure to count electron domains before predicting shape with VSEPR. A double bond counts as one direction of electron density. Separate electron-domain geometry from molecular shape, because lone pairs also occupy domains. Molecular polarity depends on the vector sum of bond dipoles and geometry.
Build your understanding
Use the detailed parent plan: two core 50–60-minute sessions per focused unit, with a third assigned for selected calculation/modeling topics. First study the core explanation and Parts 1–2 with their figures/checks; next study Parts 3–4, the whole-unit application, worked cases, varied practice and exit check. Attempt reasoning before opening guidance. Extend time when actual understanding requires it; investigations and chapter assessments have separate sessions.
Part 1 · Count valence electrons including charge
A Lewis structure tracks valence-electron pairs and atom connectivity. Add the valence electrons from all atoms, then add electrons for negative charge or subtract them for positive charge. In CO₂ the total is sixteen: eight in two double bonds and eight as oxygen lone pairs. A line represents two electrons, while a dot represents one. Counting the whole drawing prevents plausible-looking structures that secretly create or omit electrons; it is a model of valence bookkeeping, not a complete electron-density map.
Read the visual model
Identify the labeled structures and units, then trace the arrows or compare values under the caption’s stated conditions. Separate the displayed evidence from a further inference. Write your reasoning before opening the guidance.
How many lone-pair electrons are shown in the CO₂ structure after its bonding electrons are counted?
Visual reasoning — after your attempt
Eight lone-pair electrons remain after eight of the sixteen are assigned to the two double bonds.
Add valence electrons from each atom, add electrons for negative charge, and subtract for positive charge. Start a reasonable skeleton using supplied connectivity, then distribute electrons while tracking the total. A Lewis structure is an electron-bookkeeping model, not a complete image of quantum electron density.
CO₃²⁻ has 4+3×6+2=24 valence electrons. Ignoring the 2− charge leaves the wrong total and can produce an inconsistent diagram.
Pause and explain
Why does a negative charge increase the valence-electron total?
Check your reasoning after answering
The ion has gained electrons relative to the neutral atom total. Each unit of negative charge adds one electron to the bookkeeping.
Part 2 · Use octets with known exceptions
The octet rule is a useful pattern for many main-group Lewis structures, especially common second-period compounds. It has important exceptions: hydrogen uses a duet, some species are electron-deficient, some have an odd electron, and heavier-element Lewis accounts can assign more than eight electrons around a center. These patterns do not mean atoms consciously seek an octet. Advanced bonding models refine the interpretation of hypervalent species rather than simply proving that empty d orbitals act as storage boxes.
Read the visual model
Identify the labeled structures and units, then trace the arrows or compare values under the caption’s stated conditions. Separate the displayed evidence from a further inference. Write your reasoning before opening the guidance.
Why is an eleven-valence-electron molecule a warning against requiring paired octets everywhere?
Visual reasoning — after your attempt
An odd total cannot be distributed exclusively into paired electron groups; an odd-electron model or a different species assignment must be considered.
For many main-group examples, octets help guide electron placement after bonding and lone pairs are assigned. Hydrogen uses a two-electron arrangement. Electron-deficient species, odd-electron species, and some heavier-element representations require exceptions. Do not force a wrong electron total merely to draw every atom with eight.
NH₄⁺ has four N–H bonds and no nitrogen lone pair in its introductory structure. The charge and hydrogen duet rule are essential to checking the total.
Pause and explain
Should an octet rule override the supplied electron count?
Check your reasoning after answering
No. Keep the correct total and recognize known exceptions. An octet is a useful pattern for many cases, not an unconditional command.
Part 3 · Use formal charge as bookkeeping
Formal charge assigns bonding electrons equally between bonded atoms: valence count minus nonbonding electrons minus half the bonding electrons. In NH₄⁺, nitrogen has four single bonds and no lone pair, giving formal charge +1; each hydrogen has zero. The sum must equal the species’ net charge. Formal charge helps compare Lewis candidates but is not a direct measurement of actual partial charge, and it is distinct from oxidation-number bookkeeping, which assigns bonding electrons differently.
Read the visual model
Identify the labeled structures and units, then trace the arrows or compare values under the caption’s stated conditions. Separate the displayed evidence from a further inference. Write your reasoning before opening the guidance.
Why must the formal charges in a complete NH₄⁺ drawing sum to +1?
Visual reasoning — after your attempt
The electron inventory represents a species with one fewer electron than the neutral-atom sum; the assignments must reproduce that total charge.
Formal charge=valence count−nonbonding electrons−half bonding electrons. The formal charges sum to the total ion charge. This assignment is not identical to measured partial charge or oxidation number. It helps compare plausible Lewis representations when combined with connectivity and chemical evidence.
An oxygen with three lone pairs and one single bond has formal charge 6−6−1=−1. An oxygen with two lone pairs and a double bond has formal charge zero.
Pause and explain
Is formal charge the actual localized physical charge on an atom?
Check your reasoning after answering
Not exactly. It is an equal-sharing bookkeeping assignment; actual electron density and partial charges can be distributed differently.
Part 4 · Interpret resonance as one delocalized structure
Resonance contributors keep the same atom connectivity while changing allowed electron assignments. In the selected ozone account, the central oxygen has +1 formal charge and one terminal oxygen −1 in either contributor. The real molecule is not alternately one localized single/double pattern and then the other; it has a delocalized electronic structure. Equivalent contributors support equivalent bond behavior in the ideal symmetric case. Resonance arrows are not ordinary chemical-reaction arrows or movement of atomic nuclei between structures.
Read the visual model
Identify the labeled structures and units, then trace the arrows or compare values under the caption’s stated conditions. Separate the displayed evidence from a further inference. Write your reasoning before opening the guidance.
What stays fixed when moving between these ozone resonance contributors?
Visual reasoning — after your attempt
The three-oxygen connectivity and nuclear arrangement stay fixed; only the formal electron and bond assignments change.
Resonance contributors have the same atom connectivity but different electron placements. The actual structure is not rapidly flipping between separate drawings. Equivalent contributors can represent delocalization and equalized bonds; bond lengths provide evidence that a single localized drawing is incomplete.
Carbonate has three equivalent placements of one drawn C=O bond. The real ion has equivalent C–O bonds in the usual model, rather than one permanently double bond and two permanently single bonds.
Pause and explain
What stays fixed between resonance contributors?
Check your reasoning after answering
Atom connectivity stays fixed while electron placement changes. They represent contributions to one delocalized structure, not different isomers.
Apply the whole unit · Account for carbonate resonance
Illustrative ion CO₃²⁻ has carbon central and three oxygen neighbors. Use C=4 and O=6 valence electrons, and the formal-charge rule supplied in this lesson.
Your task
Count electrons, draw a valid contributor with one double bond, assign formal charges, draw the other equivalent contributors, and explain what equal measured C–O lengths would support.
What your response should show
- Choose the appropriate chemical model or relationship and state its conditions.
- Show substitutions, units, particle or charge counts, and intermediate reasoning as relevant.
- Use supplied evidence to justify the conclusion; distinguish observation from inference.
- Identify a meaningful limitation, uncertainty, or next measurement.
Response guidance — after your attempt
Count 24 valence electrons.
Use one C=O and two C–O bonds; assign lone pairs to satisfy oxygen octets.
Carbon and double-bonded oxygen have formal charge 0; each single-bonded oxygen has −1.
Move the drawn double bond to each oxygen while keeping connectivity unchanged.
The formal charges total −2 and all 24 electrons are accounted for. Equal C–O lengths support delocalization rather than three permanently unequal localized bonds. Contributors are bookkeeping representations of one ion; they are not atom rearrangements or a movie of the ion switching structures.
Scientific references and source-model foundations
Parent / teacher teaching notes
Check total electrons before judging a diagram. Separate formal charge, oxidation number, and partial charge. Use measured-bond evidence to explain why resonance is needed. Ask the learner to explain each particle symbol, axis, unit and arrow before calculating. Distinguish hypothetical models and calculated curves from measured records. Public OpenSciEd excerpts here are selected foundation prompts, not complete high-school teacher editions or an endorsement. All required advanced explanations and practice are provided locally.
The key idea
Count electrons and charge before predicting bonding, shape, and polarity.
Worked examples
Example 1
- O contributes 6 electrons and two H contribute 2: total 8.
- Two O–H bonds use 4 electrons.
- The remaining 4 give two oxygen lone pairs.
Example 2 · Worked application
- Carbon supplies 4 and two oxygens supply 12: total 16.
- Place carbon centrally and form two double bonds.
- Each oxygen has two lone pairs; count all 16 electrons and check octets.
Formal electron counting is a model; resonance is not rapid physical switching.
Example 3
- C contributes 4 and three O contribute 18.
- Add 2 for the negative charge.
Example 4 · Find and fix the mistake
- Test the claim against the stated conditions: Formal electron counting is a model; resonance is not rapid physical switching.
- Resonance drawings represent contributions to one delocalized electronic structure, not distinct structures flipping back and forth.
Formal electron counting is a model; resonance is not rapid physical switching.
Example 5 · Reason from the visual model
- Formal charge assigns bonding electrons equally between bonded atoms: valence count minus nonbonding electrons minus half the bonding electrons. In NH₄⁺, nitrogen has four single bonds and no lone pair, giving formal charge +1; each hydrogen has zero. The sum must equal the species’ net charge. Formal charge helps compare Lewis candidates but is not a direct measurement of actual partial charge, and it is distinct from oxidation-number bookkeeping, which assigns bonding electrons differently.
- The electron inventory represents a species with one fewer electron than the neutral-atom sum; the assignments must reproduce that total charge.
Original ammonium Lewis bookkeeping. The flat cross shows connectivity only, not the molecule’s tetrahedral geometry.
Avoid these mistakes
- Formal charge is bookkeeping; resonance drawings are not separate alternating molecules.
Check your understanding
How do you check a proposed Lewis structure?
Discuss the answer
Recount valence electrons, confirm total charge, check bond and lone pairs and applicable octet exceptions, then determine shape.
Return to the phenomenon
Return to the initial case. Do two resonance drawings show a molecule changing back and forth? Revise your drawing or explanation using two named ideas and one stated limitation.
Compare explanations after your attempt
They represent one delocalized structure, not a time sequence of two isolated structures. The model should explain equivalent bond behavior and conserve electron count.
What should we investigate next?
Which comparison would test a remaining uncertainty? State what you would change, keep comparable and measure.
Investigate, calculate and revise
Keep a first model and an evidence-supported revision. Attempt each task before opening guidance; all required explanations and supplied-data routes are available locally.
Activity 1 · Return to the opening case
Use the opening phenomenon, unit definitions and explained Parts 1–4.
Your task
Return to the initial case. Do two resonance drawings show a molecule changing back and forth? Revise your drawing or explanation using two named ideas and one stated limitation.
Success criteria
- Use the stated observations and appropriate units.
- Connect a labeled model with the chemical mechanism.
- State assumptions and an evidence limitation.
- Explain a substantive revision or next test.
Hint
Keep species identity, material transfer and energy account distinct.
Response guidance — discuss after your attempt
They represent one delocalized structure, not a time sequence of two isolated structures. The model should explain equivalent bond behavior and conserve electron count.
Activity 2 · Transfer to a changed case
Cover the worked answer and use the explained models, relationships and notebook.
Your task
For NH₄⁺, count valence electrons, assign nitrogen’s formal charge for four N–H single bonds and no lone pair, and distinguish that charge from molecular shape.
Success criteria
- Use the stated measurements or conditions, with correct units.
- Connect the result to the chemical mechanism or evidence relationship.
- Identify the assumption, model limit or further evidence requested.
Hint
Name known quantities and conditions, show intermediate steps and check conservation.
Response guidance — discuss after your attempt
Total valence electrons are 5+4(1)−1=8. Nitrogen formal charge is 5−0−4=+1 and each H is zero. The flat Lewis connectivity does not specify measured angles; the common molecular arrangement is tetrahedral.
Words to know
- lone pair
- A valence electron pair not represented as a bond.
- formal charge
- An electron-counting assignment, distinct from partial charge.
- resonance
- Multiple electron drawings representing delocalized bonding in one structure.
In the real world
Use the chapter investigation to connect this unit to measured evidence and a practical decision.
Exit check
How do you check a proposed Lewis structure?
Check your explanation after answering
Recount valence electrons, confirm total charge, check bond and lone pairs and applicable octet exceptions, then determine shape.
Explain your reasoning without notes. If a key idea is missing, revisit the lesson and try again.