Module 4 Notes: Lewis Structures, Shapes, and Intermolecular Forces

Module focus: Lewis structures, shapes of molecules, and intermolecular forces

Key questions:

• What are Lewis structures and how to construct them?
• What is electronegativity?
• What is bond polarity?
• How do Lewis structures determine 3D shapes (molecular geometries) via VSEPR?
• What are intermolecular forces and their origins?

Lewis structures: basic idea

• Represent covalent bonding by showing valence electrons in a molecule
• Bonding pair = two electrons shared between two different atoms; represented by a single line
• Lone pair = two electrons localized on a single atom; represented by two dots
• Example: a molecule with one bonding pair and three lone pairs around the central atom

Skeleton structure (Lewis skeleton)

• Central atom: the atom in the middle bonded to two or more atoms
• Terminal atoms: bonded only to the central atom
• Skeleton structure shows which atoms are bonded to each other (central and surrounding atoms)

Step-by-step procedure to draw Lewis structures (the five steps)
1) Draw the skeleton structure (central atom + bonded atoms)
2) Sum the valence electrons for all atoms (from the periodic table) and write the total

• Valence electrons come from group numbers (e.g., 3A = 3, 4A = 4, etc.)
• If there is a formal charge, include extra electrons accordingly (e.g., CN⁻ adds one extra electron)
• Example: for CN⁻ total valence electrons = $4 + 5 + 1 = 10$
  3) Subtract two electrons for each bond in the skeleton (each bond uses 2 electrons)
• This accounts for the electrons already used in bonding
• Example: for H2CO (formaldehyde), there are 3 bonds from C to H/H/O → subtract $3 imes 2 = 6$ electrons
  4) Count how many electrons are needed to satisfy the octet on each atom (except H, Be, B which have special cases)
• Oxygen typically needs 6 more electrons (to reach 8 total around O in the final structure)
• Carbon typically needs 2 more electrons to reach 8
• Hydrogen can only have 2 electrons (duet rule)
• Compare electrons available after step 3 to the electrons needed in step 4
  5) If extra electrons are needed to satisfy octets, form additional bonds (each extra bond adds 2 electrons)
• If there are more electrons than needed, place remaining electrons as lone pairs on atoms to complete octets
• If you don’t have enough electrons to complete octets, add more bonds (usually single bonds) as needed

Worked example: Lewis structure for formaldehyde, H₂CO

• Central atom: Carbon; skeleton: H–C–O with C in the center and H on two sides and O attached
• Step 2: Valence electrons = $4 ext{ (C)} + 6 ext{ (O)} + 1 ext{ each for two H} = 12$
• Step 3: Bonds: three bonds from C to H, H, O → electrons used = $3 imes 2 = 6$; remaining = $12 - 6 = 6$
• Step 4: Octet needs: O needs 6 more, C needs 2 more, H needs 0 more (H has duet already)
• Total needed to complete octets = $6 + 2 = 8$; not enough electrons (only 6 left)
• Step 5: Add one more bond between C and O (double bond) to provide the extra 2 electrons
• Now carbon has 8 electrons, oxygen has 6 electrons in lone pairs plus 2 in bond (total 8)
• Final check: total electrons used = 12 (initial) and octets satisfied for all atoms except H (which is allowed to have 2)
• Resulting Lewis structure: C=O with two C–H single bonds and one C–O double bond; all atoms obey octet except H

Important point: formal charge and multiple bonds

• When necessary, form double or triple bonds to satisfy octets
• Hydrogen cannot exceed 2 electrons
• Be and B have special rules: Be may have only 4 electrons around it (exception)

Worked example: CN⁻ (cyanide anion)

• Central rule: negative charge adds electrons to the total valence electron count
• Carbon (4 valence) + Nitrogen (5 valence) + 1 extra electron from the negative charge → total = $10$ valence electrons
• Construction leads to the CN⁻ Lewis structure with a triple bond or resonance forms, consistent with the negative charge

Additional examples to practice Lewis structures

• SO₂: central S with two O atoms; total valence electrons = $18$; adjust to satisfy octets with bonding and lone pairs; final structure includes correct distribution of lone pairs and bond orders
• NCl₂F: central N with two Cl and one F; total valence electrons = $26$; form skeleton; distribute electrons to satisfy octets; ensure valence electrons are used up exactly
• BeCl₂: Be is central; SN = 2 (see below); no lone pairs around Be; two Cl attached; perfect linear arrangement
• HCN: central C with H and N; triple bond between C and N; SN = 2; linear geometry
• BF₃: central B with three Cl; no lone pairs; SN = 3; electron pair geometry = trigonal planar (120° angles)
• CH₄: central C with four H; SN = 4; electron pair geometry = tetrahedral (109.5° angles)
• NH₃ (ammonia): SN = 4; one lone pair; electron pair geometry = tetrahedral; molecular geometry = trigonal pyramidal
• H₂O (water): SN = 4; two lone pairs; electron pair geometry = tetrahedral; molecular geometry = bent (V-shaped) after removing lone pairs

Steric number (SN) and electron pair versus molecular geometry

• Steric number = number of lone pairs on the central atom + number of atoms bonded to the central atom
• SN determines the electron pair geometry (the arrangement of electron domains around the central atom)
• Common SN and electron pair geometries:
• SN = 2 → electron pair geometry: linear (bonded pairs: 2; angle ≈ 180°)
• SN = 3 → electron pair geometry: trigonal planar (three electron domains; angles ≈ 120°)
• SN = 4 → electron pair geometry: tetrahedral (four electron domains; angles ≈ 109.5°)

Examples of electron pair geometry vs molecular geometry

• BeCl₂: SN = 2; electron pair geometry = linear; molecular geometry = linear (no lone pairs on Be)
• HCN: SN = 2; linear electron geometry; molecular geometry = linear
• BF₃: SN = 3; electron pair geometry = trigonal planar; molecular geometry remains trigonal planar
• CH₄: SN = 4; electron pair geometry = tetrahedral; molecular geometry = tetrahedral
• NH₃: SN = 4; electron pair geometry = tetrahedral; molecular geometry = trigonal pyramidal (one lone pair)
• H₂O: SN = 4; electron pair geometry = tetrahedral; molecular geometry = bent (two lone pairs)

Polarity: dipole moments and electronegativity

• Dipole moment: a measure of the unequal sharing of electrons in a bond (polar bond) creating a partial negative charge on the more electronegative atom and a partial positive charge on the less electronegative atom
• Polar bond vs nonpolar bond:
• If two atoms are identical (e.g., I–I in I₂), the bond is nonpolar (equal sharing)
• If two atoms are different, the bond is polar (difference in electronegativity) with a dipole pointing toward the more electronegative atom
• Electronegativities: trend across the periodic table and down a group
• Across a period (left to right): electronegativity increases
• Down a group: electronegativity decreases
• Example values (to illustrate trend): Cl ≈ $3.0$, F ≈ $4.0$, O ≈ $3.5$, P ≈ $2.1$
• Determining polarity from electronegativity differences
• Chlorine–fluorine (Cl–F) difference = $|4.0 - 3.0| = 1.0$ → polar bond
• Oxygen–fluorine (O–F) difference = $|4.0 - 3.5| = 0.5$ → polar but less than Cl–F
• Phosphorus–fluorine (P–F) difference = $|4.0 - 2.1| = 1.9$ → very polar bond
• Molecular polarity depends on both bond polarity and molecular geometry (dipole moments may cancel in symmetric molecules)

Visualizing dipoles and polarity (conceptual)

• Dipole arrows point toward the more electronegative atom
• Arrows may cancel in symmetric geometries (nonpolar molecules like CO₂ or CCl₄)
• If arrows do not cancel, the molecule has a net dipole moment (polar molecule)
• Tug-of-war analogy: two opposite dipoles with equal strength cancel to yield nonpolar; unequal strengths lead to a net dipole

Intermolecular forces (IMFs) and their impact on states of matter

• Intermolecular forces are attractions or repulsions between molecules; they influence whether a substance is a gas, liquid, or solid
• Energy perspective: solid → attraction energy >> kinetic energy; liquid → attraction energy ≈ kinetic energy; gas → attraction energy << kinetic energy
• Three main types of IMFs discussed:
• Dipole–dipole attractions (between polar molecules; strength depends on dipole moments)
• London dispersion forces (present in all molecules; especially stronger for larger, more polarizable molecules; weak in small molecules)
• Hydrogen bonding (strongest among the three; occurs when H is bonded to N, O, or F and attracted to a lone pair on another N, O, or F)

London dispersion forces (LDF)

• Origin: instantaneous dipoles induced by momentary uneven electron distribution; these induce dipoles in neighboring molecules
• Magnitude increases with molecular size (polarizability) and surface area; larger molecules have stronger LDF
• Even nonpolar molecules experience LDF; order of magnitude of boiling points correlates with LDF strength and molar mass
• Examples illustrating LDF vs size:
• CH₄ (gas, small) vs SNH₄ (larger) have different boiling points due to dispersion forces
• Halogens: F₂, Cl₂, Br₂, I₂ show increasing boiling points with increasing molecular mass, largely due to dispersion forces; I₂ is solid at room temperature due to very strong dispersion forces
• LDF can dominate in very large nonpolar molecules, sometimes surpassing dipole–dipole contributions in importance for boiling points

Hydrogen bonding

• A specific type of strong IMF: H attached to N, O, or F in one molecule can form a hydrogen bond with a lone pair on N, O, or F of another molecule
• Requires: (i) a hydrogen atom bound to N, O, or F, and (ii) a lone pair on the heteroatom of another molecule
• Examples and implications:
• Water (H₂O): strong hydrogen bonding explains high boiling point relative to other small molecules and why ice is less dense than liquid water (ice expands because of structured hydrogen-bonding network)
• Ammonia (NH₃), HF, NH₃, H₂O, and others exhibit hydrogen bonding to varying extents

Polarity and IMFs: applying to molecular predictions

• Nonpolar molecules (no net dipole) tend to rely on London dispersion forces (e.g., BF₃, Br₂, O₂, CO₂ in certain contexts)
• Polar molecules can have dipole–dipole interactions in addition to dispersion forces; hydrogen bonding can provide extra strength when applicable
• Example comparisons:
• BF₃ vs BrF₃: both trigonal planar; BF₃ is nonpolar (arrows cancel), BrF₃ has stronger interactions due to heavier atoms and potential dipole contributions
• NHCl₂ vs NBr₃: both polar, but NHCl₂ can engage in hydrogen bonding (N–H) whereas NBr₃ cannot; NHCl₂ tends to have the stronger overall intermolecular interactions due to hydrogen bonding

Quick practice problems (conceptual solutions)

• Is H₂S polar or nonpolar?
• H₂S has SN = 4 (two lone pairs on S, two bonded H atoms); electron pair geometry = tetrahedral; molecular geometry = bent; dipole moment remains (polar) → H₂S is polar
• Is SiCl₄ polar or nonpolar?
• SiCl₄ has SN = 4; electron pair geometry = tetrahedral; all substituents are the same (Cl); polar bonds cancel → SiCl₄ is nonpolar

Recap: connecting concepts across the module

• Lewis structures are foundational to predicting shapes and polarity
• VSEPR (Valence Shell Electron Pair Repulsion) model uses the steric number to predict electron-pair geometry
• Molecular geometry is derived by removing lone pairs from the electron-pair geometry to focus on the arrangement of atoms
• Bond polarity arises from electronegativity differences; dipole moments indicate molecular polarity when not canceled by geometry
• IMFs determine phase behavior and properties; dipole–dipole interactions, London dispersion forces, and hydrogen bonding govern how strongly molecules attract one another
• Real-world relevance: boiling points, phase transitions, ice density, and solubility are tied to the strengths and types of IMFs

Final practice synthesis example: comparing intermolecular forces in two pairs

• Pair 1: BF₃ vs BrF₃ (both polar/nonpolar depending on geometry and substituents)
• BF₃: polar bonds but symmetric trigonal planar geometry → nonpolar molecule overall; dispersion is present but weaker due to smaller size
• BrF₃: polar bonds and less symmetry; overall dipole moment exists; stronger attractions due to larger size and dipole contributions
• Pair 2: NHCl₂ vs NBr₃
• Both polar; both have dipole–dipole forces
• NHCl₂ has N–H bond enabling hydrogen bonding in appropriate conditions; NBr₃ lacks N–H or O–H or F–H bonds for hydrogen bonding; thus NHCl₂ typically stronger overall due to hydrogen bonding capability

Practical takeaway for exam-style questions

• Be able to draw the Lewis structure for a given molecule or ion, count valence electrons, and determine the correct bonding pattern (single, double, triple bonds) to satisfy octets (or duet for H)
• Identify the central atom, calculate the steric number, and state both the electron-pair geometry and molecular geometry
• Assess bond polarity and overall molecular polarity by considering electronegativity differences and geometry
• Describe the three main IMFs and recognize which ones are present in a given molecule; connect IMFs to observed properties like boiling point and density (e.g., ice density, boiling points across a group)

Quick reference formulas and numbers (for quick recall during problems)

• Bonding electrons per bond: $2$
• Octet rule: atoms (except H) prefer $8$ electrons around them; hydrogen prefers a duet of $2$ electrons
• Steric number (SN) = lone pairs on central atom + atoms bonded to central atom
• Typical electron-pair geometries: linear (SN=2, 180°), trigonal planar (SN=3, 120°), tetrahedral (SN=4, 109.5°)
• Important angles: $180^{<br /> ing}$, $120^{<br /> ing}$, $109.5^{<br /> ing}$
• Dipole moment concept: polar vs nonpolar depending on whether bond dipoles cancel; arrows indicate direction toward more electronegative atom; magnitude relates to Δelectronegativity
• Intermolecular forces: Dipole–dipole, London dispersion forces (LDF), Hydrogen bonding
• Trends in electronegativity: across a period increases; down a group decreases
• Hydrogen bonding condition: H–N, H–O, or H–F with a lone pair on the other molecule’s N, O, or F; explains water’s unique properties such as ice density

Connections to broader chemistry foundations

• Builds on periodic trends and electron configuration learned in earlier modules
• Demonstrates how microscopic electron distributions influence macroscopic properties (boiling points, phases, densities)
• Provides a framework for predicting molecular behavior in real-world contexts (solubility, reactivity, and material properties)