Sep 17: Chapter 4 & 5: Lewis Structures, Resonance, and VSEPR — Comprehensive Notes

Lecture Notes: Lewis Structures, Resonance, and Molecular Geometry (Chapters 4–5)

• Recap of in-class plan and study tips

  • Today’s recitation focus: naming molecules and writing molecular formulas from names; practice problems on naming; in-recitation help if students struggle with naming

  • In-class worksheet on Lewis structures to sharpen skills; name molecules, build Lewis structures, and write formulas from names

  • End of Chapter 4 today; Chapter 5 introduces molecular geometry (3D shapes) and will be covered Monday; a test (Test 1 material) likely Monday, with a review day Thursday for any remaining questions

  • We are still in Chapter 4 (Lewis structures) and comparing ionic vs covalent bonding

  • Ionic Lewis structures: electrons are transferred (ionic transfer), not what we focus on for covalent bonds

• Core focus of Chapter 4: Covalent Lewis structures

  • Goal: identify electron pairs, use them to form bonds, and determine where bonding pairs live

  • Four rules for covalent Lewis structures (recap)

  1. Calculate valence electrons

  2. Set up a skeletal structure with the center atom having the highest bonding capacity (electronegativity and unpaired electrons influence this)

  3. Convert electron dots to bonding pairs (dots → lines representing bonds)

  4. If any atom lacks an octet, fill it by creating double or triple bonds via sharing more electrons

• Key concepts and definitions

  • Valence electrons: electrons in the outermost shell of an atom (e.g., H = 1, C = 4, O = 6, N = 5, etc.)

  • Bonding capacity (bonding ability) relates to the number of unpaired electrons and electronegativity

  • Octet rule: atoms tend to complete an octet (8 electrons) around them in Lewis structures, though there are notable exceptions

  • Electronegativity (EN): tendency of an atom to attract electrons; higher EN means stronger attraction

  • EN trends: EN increases as you go up and to the right on the periodic table

  • Lone pairs: nonbonding electron pairs on atoms in Lewis structures

  • Bonding pairs: pairs of electrons shared between atoms forming a bond

• Example 1: Acetylene, C₂H₂ (molecular formula)

  • Total valence electrons: $V = 2 imes 4 + 2 imes 1 = 10$

  • Skeleton setup: center atoms chosen by highest bonding capacity; for acetylene, the structure is H–C≡C–H with C–C bond in the middle

  • Initial representation (based on valence): one C–C bond and two C–H bonds placed, giving a total electron count that must satisfy octets

  • Octet completion: convert lone pairs into bonding pairs to reach complete octets for both carbons

  • Final Lewis structure: H–C≡C–H (each carbon forms four bonds total: one to hydrogen and three to the other carbon via a triple bond)

  • Important notes:

  • The structure satisfies octets for both carbons (8 electrons around each C)

  • This example illustrates converting lone pairs to bonding pairs to achieve octets

• Example 2: Ozone, O₃ (O₃) and resonance

  • Concept recap: ozone shows resonance due to a symmetric tri-atomic arrangement with equivalent atoms

  • All three oxygens share the same element; multiple valid Lewis structures exist (resonance forms)

  • Total valence electrons: for O₃, $V = 3 imes 6 = 18$

  • Draw a skeletal O–O–O with central oxygen and place electrons to satisfy octets

  • Convert lone pairs to bonding pairs to satisfy octets; distribute electrons so that each oxygen reaches an octet

  • Resonance forms arise when a molecule with multiple valid Lewis structures exists; the real structure is the average (hybrid) of all resonance forms

  • Definition: Resonance structure is when a single Lewis structure cannot represent the actual delocalized bonding; the molecule is described by a resonance hybrid

  • Example of resonance in O₃:

  • Structure A: O=O–O (double bond on left) and single bond to right O with lone pairs adjusted

  • Structure B: O–O=O (double bond on right) and single bond to left O with lone pairs adjusted

  • Real structure: average of Structures A and B (resonance hybrid)

  • Concept note: The resonance forms are not separate real molecules; they are contributing forms that collectively describe electron distribution

• Example 3: Nitrate ion, NO₃⁻ (polyatomic ion)

  • Valence count

  • Nitrogen: 5 valence electrons

  • Oxygen: 6 valence electrons each, 3 O → 3×6 = 18

  • Negative charge adds 1 electron

  • Total valence electrons: $V = 5 + 18 + 1 = 24$

  • Skeleton: place nitrogen in the center with three oxygens around it (N in the middle is chosen because it typically has the highest unpaired electron count among the atoms involved)

  • Fill octets: place bonding electrons to satisfy octets on all atoms; adjust as needed to satisfy the octet around nitrogen and oxygens

  • Formal charge approach (to determine the most stable resonance form):

  • Definition: formal charge (FC) for an atom A is
    FC(A) = VA - ig(N{ ext{nonbonding}}(A) + frac{1}{2}N_{ ext{bonding}}(A)ig)
    (Note: in this lecture, the calculation is described as using nonbonding electrons plus bonding electrons; some treatments use half of the bonding electrons; both approaches lead to the same result when applied consistently to all atoms in a structure.)

  • Total FC sum must equal the overall charge of the species (0 for neutral molecules; equal to the ion charge for ions)

  • Stability rules for formal charges (guides):

    • Ideally, all atoms have FC = 0

    • If not possible, more stable structures have the fewest atoms with nonzero FC

    • If nonzero FC exists, place negative FC on the most electronegative atom (electronegativity trend: higher EN toward the top-right of the periodic table)

  • Example NO₃⁻ resonance forms:

    • There are three equivalent resonance structures with one N–O bond being a single bond and the other two being double bonds, distributed so that the overall -1 charge resides on the oxygens via resonance

    • In each form, the formal charges balance to give total charge −1; the average yields the resonance hybrid

  • Practical notes:

  • For nitrate NO₃⁻, three equivalent resonance forms exist; the negative charge is delocalized over the three oxygens

  • The most electronegative atom (oxygen) bears the negative charge in the resonance forms that distribute nonzero FCs

• Example 4: Dinitrogen monoxide, N₂O (nitrous oxide)

  • Valence count: N (5) + N (5) + O (6) = 16 valence electrons

  • Skeleton: N–N–O with the center atom chosen to maximize the number of unpaired electrons (often the terminal N can be placed in the center depending on the bonding scenario; here the lecturer places N in the center)

  • Build Lewis structures to complete octets; consider possible resonance forms

  • Formal charge analysis:

  • Calculate FC for each possible resonance form

  • The total FC sum across a structure must be zero for a neutral compound; for ions, the sum equals the charge

  • Among possible resonance forms, the most stable form is typically the one with the negative FC on the most electronegative atom (oxygen) when a nonzero FC exists

  • Takeaway: N₂O has multiple resonance forms; the most stable form is the one where negative charge resides on oxygen (the most electronegative element), and the overall charge distribution is consistent with the molecule's neutrality

  • Note: In the lecture, the instructor shows a few resonance forms and demonstrates how formal charges guide the selection of the most stable resonance structure

• Example 5: Carbon dioxide, CO₂

  • Valence count: C (4) + 2×O (2×6) = 4 + 12 = 16 valence electrons

  • Skeleton: O=C=O (central carbon double-bonded to each oxygen)

  • Fill octets: carbon forms two double bonds to oxygen; each O has an octet

  • Formal charges: all atoms have FC = 0 in the dominant resonance form; this is the most stable arrangement

  • Resonance forms: CO₂ has equivalent resonance forms with two C=O bonds; the actual structure is a resonance hybrid with equal bond character between C and O

• Summary: Formal charges, resonance, and octet rules in Chapter 4

  • Formal charge helps determine the most stable resonance form, especially when multiple valid Lewis structures exist

  • Resonance forms can be equivalent (leading to a resonance hybrid where the actual bonds are delocalized) or non-equivalent (one form may be more stable due to FC distribution)

  • When calculating FCs, ensure the sum of all FCs matches the overall charge of the molecule/ion

  • The stability preference often described as:

  • Zero FC on all atoms is ideal but not always possible

  • If nonzero FC exists, place negative FC on the most electronegative atoms

• Transition to Chapter 5: Molecular geometry and VSEPR (valence shell electron pair repulsion)

  • Chapter 5 builds on Lewis structures to understand 3D shapes of molecules

  • VSEPR principle: electron pairs (bonding and lone pairs) around the central atom arrange themselves to minimize repulsion, thereby minimizing molecular energy

  • Key definitions

  • Electron pair geometry: spatial arrangement of all electron pairs (bonding and lone pairs) around the central atom

  • Molecular geometry: actual arrangement of atoms in space (shape of the molecule) based on electron pair geometry

  • If there are no lone pairs on the central atom, electron pair geometry equals molecular geometry

• How to determine geometry: Steric number (SN)

  • Definition: $ext{SN} = n<em>{ ext{bonded}} + n</em>{ ext{lone}}$

  • SN values and corresponding geometries (the five common shapes)

  • SN = 2 → Linear geometry (bond angle ≈ 180°)

  • SN = 3 → Trigonal planar geometry (bond angle ≈ 120°)

  • SN = 4 → Tetrahedral geometry (bond angle ≈ 109.5°)

  • SN = 5 → Trigonal bipyramidal geometry (bond angles: equatorial ~120°, axial ~90°)

  • SN = 6 → Octahedral geometry (bond angles ~90°)

  • As SN increases, bond angles generally decrease due to crowding, causing the structure to “collapse” toward smaller angles

• Examples and notes for SN-based geometries

  • Carbon dioxide, CO₂: central carbon with two bonded atoms (O) and no lone pairs on C

  • SN = 2, Linear geometry, bond angle ≈ 180°

  • Boron trifluoride, BF₃: central boron with three bonds and no lone pairs

  • SN = 3, Trigonal planar geometry, bond angles ≈ 120°

  • Important exception to the octet rule: boron in BF₃ does not complete an octet (only six electrons around B), illustrating an electron-deficient compound

  • Methane, CH₄: central carbon with four bonds and no lone pairs

  • SN = 4, Tetrahedral geometry, bond angles ≈ 109.5°

  • Phosphorus pentachloride, PCl₅: SN = 5, Trigonal bipyramidal geometry (two distinct sets of bond angles: 120° in the plane and 90° perpendicular to it)

  • Sulfur hexafluoride, SF₆: SN = 6, Octahedral geometry, bond angles ≈ 90°

• Practical notes and tips from the lecture

  • The instructor emphasizes that rules and numbers are guides to help orient thinking; problems on tests will often require applying several steps in sequence (valence calculation, skeleton, octet completion, resonance, formal charges) rather than rote memorization

  • When working with resonance and formal charges, start by drawing all plausible Lewis structures, then compute FCs for each atom in each structure

  • The stability criterion for resonance forms often requires placing the negative FC on the most electronegative atom

  • Equivalence matters: if resonance structures are equivalent, their FCs are the same and the real structure is the average of the contributing forms

  • With non-equivalent resonance structures, use FC to select the most stable (the one with the most zero FCs and the negative FC on the most EN atom)

  • The lesson also includes reminders of practical test-taking strategies: practice naming and deriving molecular formulas from names; Lewis structure worksheets will reinforce this skill and prepare students for in-class questions

• Quick reference formulas and concepts (summary)

• Valence counting and electron total

  • Total valence electrons for a molecule: $V = ext{sum of valence electrons of all atoms involved}$

  • Example counts:

  • Acetylene: $V = 2 imes 4 + 2 imes 1 = 10$

  • Ozone: $V = 3 imes 6 = 18$

  • Nitrate NO₃⁻: $V = 5 + 3 imes 6 + 1 = 24$

  • Dinitrogen monoxide N₂O: $V = 2 imes 5 + 6 = 16$

  • Carbon dioxide CO₂: $V = 4 + 2 imes 6 = 16$

• Formal charge calculation (FC)

  • For atom A: FC(A) = VA - ig(N{ ext{nonbonding}}(A) + N_{ ext{bonding}}(A)ig)

  • Total FC in a neutral molecule sums to 0; total FC in an ion sums to the charge of the ion

  • Stability rules:

  • Ideally, FC = 0 for all atoms

  • If nonzero FC occurs, place negative FC on the most electronegative atom

• Resonance basics

  • Resonance forms are alternative Lewis structures for the same molecule; the actual structure is a hybrid of these forms

  • Equivalent resonance forms have the same FC on corresponding atoms and contribute equally to the hybrid

  • Non-equivalent resonance forms require FC analysis to identify the most stable contribution

• VSEPR and geometry workflow

  • Build Lewis structure first, then compute the steric number (SN)

  • Determine geometry from SN

  • Use geometry to predict bond angles and spatial arrangement

• Practical exam-oriented reminders

  • Expect problems that ask you to: draw Lewis structures from names or formulas, calculate valence electrons, determine formal charges, identify the most stable resonance form, and predict molecular geometry using VSEPR

  • Practice with these core examples (acetylene, ozone, NO₃⁻, N₂O, CO₂) and with simple geometry cases (CO₂, BF₃, CH₄, PF₅, SF₆) to gain fluency

• Sidebar: a life science-side aside mentioned in the lecture

  • The ozone layer is an example of an allotrope issue: same element (oxygen) with different molecular forms (O₃ versus O₂)

  • Allotropes are different structural forms of the same element; ozone plays a role in absorbing ultraviolet light and has complex environmental effects

  • This aside is not typically tested in depth for the Lewis structure unit, but it helps connect chemistry to real-world context

• Closing notes from the instructor

  • A worksheet on Lewis structures is planned for tomorrow to reinforce valence calculations, resonance, and formal charges

  • Chapter 5 will build on this, focusing on 3D shapes and geometry (VSEPR)

  • The class will likely have a Monday test covering these topics, with a Thursday review for any remaining questions

  • The instructor also notes personal difficulty with quick mental math and emphasizes using structured notes and stepwise reasoning to prevent mistakes