Chapter 1-6 Notes: Acidity, Electron Density, and Proton Transfer (Vocabulary Flashcards)

Electronegativity and acidity trends

• Trend relates to electronegativity: as we decrease the electronegativity of the atom that the hydrogen is attached to, that particular hydrogen becomes more acidic.
  • Example framing in the transcript: nitrogen must have three (ammonia, NH₃); carbon has four (methane, CH₄). As we go across the series, one proton among the hydrogens can be more acidic.
• As we increase the electronegativity on the atom attached to H, the bond becomes more polarized. The electrons in the bond spend more time on the electronegative atom, weakening the O–H/N–H/C–H bond and making proton loss easier.
• Result: more polarized bonds lead to easier proton release (stronger acidity) in that context.

Conjugate base formation and stabilization

• When a proton is donated by an acid (HA) to a base (B⁻), a conjugate base (A⁻) is formed and the electrons shift back toward the atom that bears the charge.
  • General reaction: HA + B^- \rightarrow A^- + HB
• The stability of the conjugate base is a key factor in acidity: more stable conjugate base → stronger acid.
• Example set across halogens (and related species):
  • Acids give conjugate bases CH₃COO⁻, NH₂⁻, OH⁻, F⁻, etc. The relative stability of these anions contributes to acidity trends.

Inductive effects (through-bond electronic effects)

• Induction refers to electronic effects that travel through sigma bonds (not through the pi system).
• Two types of inductive groups:
  • Electron-withdrawing groups (EWG) pull electron density toward themselves and away from the acidic proton, increasing acidity.
  • Electron-donating groups (EDG) push electron density toward the acidic proton, decreasing acidity.
• Key factors:
  • Proximity: how close the electronegative atom is to the hydrogen (or to the nearest neighbor potential site).
  • Number of electronegative atoms: more withdrawing atoms generally increase acidity.
• Example idea: If a carbon bears more fluorines (or chlorines) nearby, it withdraws more electron density from the adjacent C–H bond than a less substituted carbon—stabilizing the resulting anion to a greater extent.
• Practical illustrations in the transcript:
  • In a structure with an oxygen adjacent to the C–H bond, oxygen withdrawal can pull electron density through the bond and stabilize the developing negative charge on the conjugate base when the proton is removed.
  • Adding a chlorine group further withdraws electron density and can stabilize the anion, increasing acidity.
• Comparative thought experiment:
  • If you replace hydrocarbon hydrogens with three chlorines near the acidic site versus only one chlorine, the more heavily chlorinated scenario stabilizes the anion more through inductive withdrawal, leading to greater acidity.

Resonance stabilization of the anion

• Resonance can stabilize the conjugate base by delocalizing the negative charge.
• Acetic acid vs ethanol (illustrative example):
  • Acetate ion (from acetic acid) has resonance stabilization across two oxygen atoms, allowing multiple resonance forms and increased stability of A⁻.
  • Ethoxide (from ethanol) has less resonance stabilization, leading to comparatively weaker acidity.
• Key point: more resonance structures for the conjugate base → greater stabilization → stronger acid.

Phenol vs alcohol and the role of resonance and solvation

• Phenol (C₆H₅OH) is more acidic than a comparable alcohol (e.g., ethanol) due to resonance stabilization of the phenoxide anion (C₆H₅O⁻).
  • Phenoxide can form several resonance structures by moving nonbonding electrons and π electrons around the ring, spreading negative charge over the ring and oxygen.
  • Transcript notes: multiple resonance forms (e.g., 4–5 forms) indicate substantial stabilization of the phenoxide anion.
• Deprotonation of phenol by a strong base (e.g., NaOH) forms sodium phenoxide:
  • Reaction: ext{C}6 ext{H}5 ext{OH} + ext{NaOH} \rightarrow ext{C}6 ext{H}5 ext{O}^- ext{Na}^+ + ext{H}_2 ext{O}
  • The resulting salt is polar, improves solubility in water, and showcases the role of counterions and solvent in stabilizing charged species.
• Solvent effects:
  • Solvents can stabilize ions via solvation (hydrogen bonding, dipole interactions).
  • Sometimes solvents stabilize cations or anions depending on the system, influencing whether deprotonation is favorable.
• Practical takeaway: phenol is deprotonated by NaOH to form sodium phenoxide salt, not by weaker bases like some other hydroxides, illustrating acidity differences and solvent/solvation effects.

Alpha vs beta substituent effects on acidity

• Substituent position relative to the acidic site matters:
  • An electron-withdrawing group (e.g., Cl) at the alpha position (nearest to the carboxyl group) has a greater impact on acidity than at the beta position.
  • Alpha-substitution (closer proximity) more effectively stabilizes the developing negative charge on deprotonation due to stronger inductive withdrawal.
• Example discussion from transcript:
  • Compare C–H acidic sites with Cl at alpha vs beta positions; the alpha-substituted case is predicted to be more acidic due to proximity of the withdrawing group.

Propionic acid substitutions and acidity predictions

• Propionic acid substitution scenarios discussed:
  • Consider substitution of chlorine at alpha position versus beta position.
  • The alpha-chloro substituted propionic acid is predicted to be more acidic than the beta-chloro substituted one due to proximity of the electron-withdrawing group to the –OH (carboxyl) group.
• Takeaway: both proximity to the acidic proton and the number of electron-withdrawing groups contribute to acidity.

Stepwise view of acidity and base strength (illustrative examples)

• Acids and bases conceptually:
  • Acids donate a proton to a base, forming a conjugate base. The ease of proton donation is governed by the stability of the conjugate base and the ability to generate it.
  • Base strength can be described in terms of its tendency to accept a proton; a stronger acid reacts more readily with bases because its conjugate base is more stable.
• Example with acetic acid and ethanol (resonance/inductive context):
  • Acetic acid’s conjugate base (acetate) is stabilized by resonance with the two oxygens and by inductive withdrawal, making acetic acid relatively stronger than ethanol.
• Comparison of resonance and inductive effects:
  • Resonance stabilization increases acidity beyond inductive effects alone.
  • Solvent effects and hydrogen bonding in the medium can further modulate acidity and stabilization of ions.

Proton transfer basics and what counts as a base

• Proton transfer can occur via various types of bases:
  • Lone-pair bases (e.g., water, hydroxide) donate electron density to accept a proton.
  • Full negative charges (anions) can also act as bases.
  • Electrons in p orbitals (as in certain bases) can participate in proton transfer or attack; p orbitals are often more reactive than sigma-bond electrons because they are higher in energy and less tightly held.
• Mechanistic note: in many cases, a proton first associates with a base’s lone pair, then the resulting conjugate base stabilizes the negative charge via resonance or solvation.

Reactions involving p orbitals and additions to double bonds

• p orbitals participate in reactions by interacting with pi bonds; initial step often involves adding a proton across a double bond, generating a carbocation on the other carbon, making it more susceptible to further reactions.
• This is illustrated conceptually in the transcript with a discussion of orbital character and reactivity of p orbitals in organic transformations.

Examples of SN1-like transformations in alcohol to alkyl halide conversion

• tert-Butyl alcohol reacting with HCl (acid-catalyzed substitution) illustrates a common acid-base/solvolysis mechanism:
  1) Protonation of the hydroxyl group: ext{(CH}3)3 ext{COH} + ext{H}^+ \rightarrow ext{(CH}3)3 ext{COH}2^+ 2) Water leaves to form a tert-butyl carbocation: ext{(CH}3)3 ext{C}^+ + ext{H}2 ext{O}
  3) Chloride attacks the carbocation to give tert-butyl chloride: ext{(CH}3)3 ext{CCl} + ext{H}_2 ext{O} ext{ (after water leaves)}
• Key idea: the leaving group (H₂O) departs after protonation, forming a stable tertiary carbocation which is then attacked by Cl⁻ to yield the alkyl chloride.
• This example ties together concepts of proton transfer, carbocation stability, and nucleophilic attack in acid-base chemistry.

Solvent and stabilization effects in acid-base chemistry

• Solvent role:
  • Solvents stabilize ions via solvation and hydrogen bonding; the solvent can influence whether a reaction proceeds to completion by stabilizing the products or reactants.
  • In the phenoxide example, sodium phenoxide formed in water is stabilized by the polar solvent (water) and by the counterion (Na⁺).
• Practical consequence: solvents can determine when deprotonation occurs and how stable the resulting ions are, affecting reaction equilibria and product distribution.

Summary takeaways (from the transcript)

• Acidity trends are explained through electronegativity, bond polarization, and conjugate-base stabilization (via resonance and inductive effects).
• Inductive effects depend on proximity and the number of electronegative substituents; alpha-position substituents have a larger impact than beta-position substituents.
• Resonance stabilization of the conjugate base (e.g., acetate in acetic acid) can dominate acidity in many cases, sometimes more so than inductive effects.
• Phenol is measurably more acidic than aliphatic alcohols due to extensive resonance stabilization of the phenoxide anion; its deprotonation by NaOH to form sodium phenoxide demonstrates both acid-base chemistry and solvation effects.
• The solvent environment and solvation play crucial roles in stabilizing ions and influencing reaction pathways.
• Understanding acidity involves combining electronegativity, inductive effects, resonance, orbital character, and solvent effects, as illustrated by examples like acetic acid vs ethanol, phenol vs alcohol, and the tert-butanol/HCl transformation.

References to practice problems and classic examples

• Compare acetic acid (CH₃COOH) and ethanol (CH₃CH₂OH) in terms of conjugate-base stability and acidity.
• Assess the acidity of phenol (C₆H₅OH) vs ethanol, using resonance stabilization of the phenoxide anion).
• Predict the relative acidity of alpha- vs beta-chlorinated propionic acids (2-chloropropanoic acid vs 3-chloropropanoic acid).
• Explain the formation of sodium phenoxide via NaOH and why phenoxide is more stabilized in water than the neutral phenol.
• Describe the SN1-like conversion of tert-butanol to tert-butyl chloride in the presence of HCl and the role of water as a leaving group.