Comprehensive Study Notes: Polar Covalent Bonds, Acids & Bases, and Intermolecular Interactions

Polar Covalent Bonds, Electronegativity, and Bond Polarity

• Most bonds are neither fully ionic nor fully covalent. Polar covalent bonds occur when bonding electrons are attracted more strongly by one atom than the other, leading to an uneven electron distribution between atoms.
• Key concept: Electronegativity (EN) drives bond polarity. Differences in EN create partial charges and directional polarity in bonds.
• Bond polarity is depicted by a partial positive charge on the less electronegative atom (δ+) and a partial negative charge on the more electronegative atom (δ−). A crossed arrow is used to indicate the direction of electron withdrawal toward the more electronegative atom in polar bonds.
• Electrostatic Potential Maps visually show charge distribution: red regions indicate electron-rich (δ−) areas; blue regions indicate electron-poor (δ+) areas.

Electronegativity and Bond Classification

• Polar Covalent Bond: ΔEN between about 0.5 and 2.0.
  • Example: C–O bond (ENC = 2.5, ENO = 3.5; ΔEN = 1.0).
• Nonpolar Covalent Bond: ΔEN < 0.5.
  • Example: C–H bond (ENC = 2.5, ENH = 2.1; ΔEN = 0.4).
• Very polar bonds (toward ionic character) occur as ΔEN grows; but most real-world bonds are intermediate.
• Polar Covalent Bond examples from lecture:
  • C–Li: ΔEN ≈ 1.5 (ENC = 2.5, ENLi = 1.0)

Inductive Effect

• Inductive effect: shifting of electrons in a σ bond in response to the electronegativity of nearby atoms.
• Influences chemical reactivity: electron donation or withdrawal transmitted through σ bonds.
• Typical trends:
  • Li, Mg → inductively donate electrons (electron-donating groups)
  • O, N → inductively withdraw electrons (electron-withdrawing groups)

Practice: Electronegativity

• Question: Which element in each pair is more electronegative?
  • a) Li vs H
  • b) B vs Br
  • c) Cl vs I
  • d) C vs H

Practice: Predict Bond Polarity with Crossed Arrows

• For bonds: H3C–Cl, H3C–NH2, H2N–H, H3C–SH, H3C–MgBr, H3C–F
• Use the crossed arrow (δ+/δ−) convention to indicate the expected polarity direction.

Molecular Polarity and Dipole Moments

• Molecular polarity results from the vector sum of all bond polarities and lone-pair contributions.
• Dipole moment (µ) is a quantitative measure of net molecular polarity.
  • Definition: oldsymbol{BC} \, \mu = Q \times r where Q is the magnitude of the charge, and r is the distance between charges.
  • In Debye units (D): $1\ \,\mathrm{D} = 3.336 \times 10^{-30} \ \mathrm{C\cdot m}$
• Example calculation:
  • If a charge of $Q = 1.60 \times 10^{-19} \ \text{C}$ is separated by $r = 100 \times 10^{-12} \ \text{m}$, then
  • $\mu = Q r = (1.60 \times 10^{-19})(100 \times 10^{-12}) = 1.60 \times 10^{-29} \ \mathrm{C\cdot m}$
  • In Debye: $\mu = \frac{1.60 \times 10^{-29}}{3.336 \times 10^{-30}} \approx 4.80\ \mathrm{D}$
• Table: Dipole moments (selected compounds)
  • NaCl: 9.00 D
  • NH3: 1.47 D
  • CH2O: 2.33 D
  • CH3NH2: 1.31 D
  • CH3Cl: 1.87 D
  • CO2: 0 D
  • H2O: 1.85 D
  • CH4: 0 D
  • CH3OH: 1.70 D
  • CH3CH3: 0 D
  • CH3CO2H: 1.70 D
  • CH3SH: 1.52 D
  • Benzene: data not provided in lecture notes
• Lone pairs contribute significantly to dipole moments by extending electron density away from nuclei, increasing charge separation.

Molecular Polarity and 3D Representations

• Three-dimensional drawings and dipole moment directions can be predicted for molecules such as:
  • H2C=CH2, CHCl3, CH2Cl2, H2C=CCl2

Formal Charges

• Formal charge (FC) is a bookkeeping device, not a statement of actual charges; it helps predict reactivity and resonance.
• General formula: $\mathrm{FC} = V - \frac{B}{2} - N$
  • V = valence electrons of the atom in the free atom
  • B = bonding electrons (counted in bonds)
  • N = nonbonding electrons on the atom
• Formal charges help explain stability and reactivity; they often align with electrostatic potential maps.
• Example: Dimethyl sulfoxide (DMSO) formal charges
  • Sulfur can bear a positive formal charge (+1)
  • Oxygen can bear a negative formal charge (−1)
• Common formal charges (typical values):
  • Carbon: often 0 or ±1 in various structures
  • Nitrogen: can be −1, 0, or +1 depending on bonding
  • Sulfur and phosphorus: can bear positive or negative charges in charged species
• Practice: Calculate formal charges for the non-hydrogen atoms in the following:
  • (a) Diazomethane, H2C=N=N:
  • (b) Acetonitrile oxide, H3C-C≡N−O:
  • (c) Methyl isocyanide, H3C−N=C:

Resonance

• Resonance: molecules can be described by two or more Lewis structures (resonance forms) that differ only in the placement of π and nonbonding electrons; the real structure is a resonance hybrid.
• Key concepts:
  • Resonance forms are imaginary; the real molecule is a single, unchanging resonance hybrid, not a rapid interconversion between forms.
  • The only differences between resonance forms are the placement of π and lone-pair electrons; atom positions do not change.
  • Resonance forms obey the octet rule for second-row elements.
  • The more resonance forms and better electron delocalization, the more stable the resonance hybrid.
• Rules for resonance (summary):
  • Rule 1: Individual resonance forms are imaginary; real structure is a resonance hybrid.
  • Rule 2: Resonance forms differ only in π or nonbonding electron placement; atoms' positions/hybridization do not change.
  • Rule 3: Resonance forms need not be equivalent; the more stable form contributes more to the hybrid.
  • Rule 4: Resonance forms obey valency and the octet rule for second-row elements.
  • Rule 5: The resonance hybrid is more stable than any individual resonance form; more forms generally lead to greater stabilization.
• Examples:
  • Acetate ion (CH3COO−): two primary resonance forms; true structure is a resonance hybrid.
  • Benzene: two dominant resonance forms with alternating single and double bonds; experimental data show all C–C bonds are equivalent due to delocalization.
  • Allyl cation/anion: allylic lone pairs and charges distributed across three-atom groups; allyl system often shows equivalent bond lengths in certain cases.
• Drawing resonance forms (2-6) involves identifying three-atom groupings with p orbitals and moving electrons via π and lone-pair interactions.

Common Resonance Patterns

• Allylic lone pairs: lone pairs adjacent to a π bond can participate in resonance (requires curved arrows: π bond → lone pair, lone pair → π bond).
• Allylic positive charge: positive charge on a carbon adjacent to a π bond; resonance forms involve shifting electrons along conjugated systems.
• Lone pair adjacent to a positive charge: lone pair donates into adjacent π bond or π system.
• A π bond between atoms of different electronegativity: movement of π and lone-pair electrons expands delocalization.
• Conjugated π bonds in a ring: continuous delocalization around a ring with unhybridized p orbitals.

Drawing Resonance Forms (Examples and Practice)

• Practice: Draw three resonance structures for carbonate ion, CO₃²⁻.
• Practice: Draw three resonance structures for the pentadienyl radical. Radical present as a dot on a carbon; move electrons to generate other resonance forms.
• Practice: Determine whether given structures are resonance forms; explain.

Acids and Bases: Brønsted–Lowry Definition

• Brønsted–Lowry Acid: a substance that donates a hydrogen ion (proton, H⁺).
• Brønsted–Lowry Base: a substance that accepts a hydrogen ion (proton, H⁺).
• In any acid–base reaction, the conjugate base of the acid and the conjugate acid of the base are formed:
  • Example reaction: $\mathrm{H-A} + \mathrm{B:} \rightarrow \mathrm{A^-} + \mathrm{BH^+}$
• Strong acids dissociate to the right in water; weak acids dissociate only slightly.

Acidity Constant (Ka) and pKa

• Ka is the equilibrium constant for the dissociation of an acid in water: $\mathrm{HA} + \mathrm{H<em>2O} \rightleftharpoons \mathrm{H</em>3O^+} + \mathrm{A^-}$
• The acidity constant is expressed as: $K<em>a = \frac{[\mathrm{H</em>3O^+}][\mathrm{A^-}]}{[\mathrm{HA}]}$
• pKa is the negative logarithm of Ka: $\mathrm{p}K<em>a = -\log K</em>a$
• Typical ranges (as given in notes): Ka values range from about $10^{15}$ for the strongest acids to about $10^{-60}$ for the weakest acids, with corresponding pKa values spanning from around
  -$-15$ to about $-60$ on a log scale (illustrative; refer to course table for exact values). In many chemistry texts, stronger acids have negative pKa values and correspondingly very large Ka values.
• pKa values in common aqueous equilibria (selected examples):
  • HCl: pKa ≈ −7.0 (Ka ≈ 10^7)
  • HNO3: pKa ≈ −1.3
  • H3O⁺ (acid in water): pKa ≈ −1.74 (reconstructed context); note: strong acids have very small pKa on the negative side.
  • CH3CO2H (acetic acid): pKa ≈ 4.76
  • H2O: pKa ≈ 15.74
  • NH4⁺: pKa ≈ 9.25 (not in lecture notes; typical value for context)
• A useful table (from lecture notes):
  • Ethanol, CH3CH2OH: pKa ≈ 16.00 → Ethoxide CH3CH2O⁻ (weaker conjugate base)
  • Water, H2O: pKa ≈ 15.74 → Hydroxide OH⁻ (conjugate base)
  • Hydrocyanic acid, HCN: pKa ≈ 9.31 → CN⁻
  • Dihydrogen phosphate, H2PO4⁻: pKa ≈ 7.21 → HPO4²⁻
  • Acetic acid, CH3CO2H: pKa ≈ 4.76 → CH3CO2⁻
  • Phosphoric acid, H3PO4: pKa ≈ 2.16 → H2PO4⁻
  • Nitric acid, HNO3: pKa ≈ −1.3 → NO3⁻
  • Hydrochloric acid, HCl: pKa ≈ −7.0 → Cl⁻

Predicting Acid–Base Reactions from pKa Values

• Key rule: proton transfer occurs from the stronger acid to the stronger base (i.e., to the conjugate base of the weaker acid).
• In an equilibrium, the product conjugate acid must be weaker (less reactive) than the starting acid, and the product conjugate base must be weaker than the starting base.
• Examples:
  • Water (pKa ≈ 15.74) vs acetylene (pKa ≈ 25): water is the stronger acid relative to acetylene, so OH⁻ will not deprotonate acetylene significantly; acetylene is a much weaker base than OH⁻.
  • Ammonia (pKa ≈ 36) vs acetone (pKa ≈ 19): acetone is the stronger acid than ammonia; the equilibrium for proton transfer from NH3 to acetone is unfavorable.
• Practice questions include ranking acidity and predicting reaction extent based on the pKa values in Table 2-3.

The Lewis Definition: Acids and Bases

• Lewis Acid: accepts an electron pair.
• Lewis Base: donates an electron pair.
• Curved-arrow formalism depicts electron-pair flow: a curved arrow from a Lewis base (donor) to a Lewis acid (acceptor).
• Examples:
  • BF3 (Lewis acid) accepting electron density from dimethyl ether (Lewis base) to form an acid–base adduct.
  • Many species such as H2O, HCl, H3CCH2OH can act as Lewis bases or acids depending on the partner.
• Curved arrows illustrate electron redistribution without necessarily involving proton transfer.
• Lewis bases frequently include oxygen- and nitrogen-containing compounds due to lone pairs of electrons.
• Some compounds can act as both Lewis acids and bases (amphoteric behavior).

Noncovalent Interactions (Intermolecular Forces)

• Major types:
  • Dipole–dipole forces: attraction or repulsion between polar molecules due to partial charges.
  • Dispersion forces (London forces): present in all molecules; arise from instantaneous dipoles from electron distribution; can be significant cumulatively.
  • Hydrogen bonds: strong dipole–dipole interaction where hydrogen is bonded to F, O, or N and interacts with a lone pair on another electronegative atom.
• Hydrogen bonding is responsible for many physical properties (e.g., high boiling point of water) and essential in biomolecules (DNA, proteins, enzymes).
• Hydrophilic: water-loving; strongly attracted to water (e.g., table sugar).
• Hydrophobic: water-fearing; not strongly attracted to water (e.g., vegetable oil).

Organic Acids and Organic Bases

• Organic acids are characterized by a positively polarized hydrogen atom (blue in electrostatic maps) and include:
  • H atom bonded to electronegative O (O–H) as in alcohols and carboxylic acids.
  • H atom on a carbon adjacent to a C=O bond (O=C–C–H) as in certain carbonyl compounds.
• Examples of organic acids:
  • Methanol (CH3OH)
  • Acetic acid (CH3CO2H)
  • Organic acids often dissociate to give conjugate bases stabilized by resonance and inductive effects.
• Organic bases commonly contain nitrogen or oxygen with lone pairs (e.g., methylamine CH3NH2, alcohols, ethers).
• Carboxylic acids (R–COOH) are abundant in biology; at cellular pH (~7.3), they exist mainly as carboxylate anions (R–CO2−).

Stability of Negative Formal Charges (A.R.I.O framework)

• A = Atom bearing the negative charge: more electronegative atoms stabilize negative charge; larger atoms spread charge better (size matters).
• R = Resonance: delocalization stabilizes the negative charge; conjugate bases with resonance are stronger acids.
• I = Inductive effects: electron-withdrawing groups withdraw electron density through σ bonds, stabilizing the negative charge.
• O = Orbital character (hybridization): orbitals with more s-character are closer to the nucleus and stabilize negative charges more effectively.
• Practical implications:
  • Ethanol vs acetic acid: acetate anion CH3COO− is stabilized by resonance, making CH3COOH a stronger acid than CH3CH2OH.
  • CF3CH2O−: inductive withdrawal by nearby F atoms stabilizes the negative charge, increasing acidity relative to ethanol.
• Key takeaway: stronger conjugate base (less stable acid) corresponds to weaker acid; stability of the conjugate base is central to acid strength.

Examples and Applications

• Ranking acidity using ARIO factors:
  • Compare acidity of various organic compounds by evaluating electronegativity, resonance, inductive effects, and hybridization.
• Applications include predicting acidity of drugs and natural products (e.g., THC, ketoprofen, propranolol) and choosing protonation sites.

Practical Notes and Real-World Relevance

• Electronegativity and bond polarity underpin reactivity in synthesis, catalysis, and biochemistry (e.g., enzyme active sites, DNA base pairing).
• Resonance explains the partial double-bond character and stability of many anions and neutral molecules; it helps justify structures that are not easily drawn by single Lewis structures.
• Understanding acid–base equilibria (pKa, Ka) is crucial for predicting reaction outcomes, stability of compounds in different environments (solvents, pH), and designing buffers.
• Lewis acid–base chemistry broadens the realm of acid–base interactions beyond proton transfers, enabling complexation and catalysis in organometallic and organic chemistry.

Quick Practice Reminders

• Be able to identify polar vs nonpolar bonds based on ΔEN and predict δ+ / δ− directions.
• Use dipole-dipole, dispersion, and hydrogen-bond concepts to rationalize physical properties and biomolecular interactions.
• Apply the resonance rules to determine resonance hybrids and stability in given structures.
• Use pKa values to predict the outcome of proton transfer reactions and to rank acid strength.
• Employ ARIO factors to rationalize the relative acidity of organic compounds and the stability of their conjugate bases.

Note on Equations and Conversions

• Dipole moment: $\mu = Q \times r$
• 1 Debye: $1\ \mathrm{D} = 3.336 \times 10^{-30} \ \mathrm{C\cdot m}$
• Relationship between pKa and Ka: $\mathrm{p}K<em>a = -\log K</em>a, \quad K<em>a = 10^{-\mathrm{p}K</em>a}$
• Formal charge: $\mathrm{FC} = V - \frac{B}{2} - N$
• Example calculation for dipole moment (given values in lecture):
  • $Q = 1.60 \times 10^{-19} \mathrm{C}, \quad r = 100 \times 10^{-12} \mathrm{m}$
  • $\mu = Qr = 1.60 \times 10^{-19} \times 100 \times 10^{-12} = 1.60 \times 10^{-29} \mathrm{C\cdot m}$
  • In Debye: $\mu = \frac{1.60 \times 10^{-29}}{3.336 \times 10^{-30}} \approx 4.80 \mathrm{D}$

Summary of Key Takeaways

• Bond polarity arises from differences in electronegativity; not all bonds are purely ionic or covalent.
• Dipole moments quantify molecular polarity and depend on both bond dipoles and molecular geometry.
• Formal charges help assess reactivity and resonance; many stable structures rely on resonance-delocalized electrons.
• Resonance hybrids explain observed bond lengths and stability beyond single Lewis structures.
• Acids and bases can be described using Brønsted–Lowry and Lewis definitions; pKa values allow prediction of proton-transfer equilibria.
• ARIO factors provide a systematic way to rationalize the strength of organic acids and bases.
• Noncovalent interactions (dipole-dipole, dispersion, hydrogen bonds) govern many physical properties and biological processes.