Chapter 2 Notes: Acids and Bases; Functional Groups (Comprehensive)

Chapter 2 Notes: Acids and Bases; Functional Groups

  • Goals for Chapter 2 (overview)

    • Identify molecular features that make compounds polar and capable of hydrogen bonding

    • Predict general trends in physical properties (boiling points, solubilities)

    • Identify acids, bases, electrophiles, nucleophiles; compare their strengths using structure, bonding, and Ka/pKa

    • Identify nucleophiles and electrophiles in Lewis acid–base reactions; use curved arrows to show electron flow

    • Identify general classes of organic compounds; functional groups underpin properties and reactivity

    • Recognize that many organic compounds can behave as acids or bases depending on environment

    • Context: acids and bases are everywhere (food, cleaners, physiology)

  • Functional groups and polarity (intro to Chapter 2)

    • Functional groups: groups of atoms that determine properties and reactivity; acidic/basic sites

    • Citric acid and ascorbic acid (vitamin C) are water-soluble acids in citrus; ascorbic acid is essential for collagen synthesis; humans cannot synthesize it; deficiency causes scurvy

  • 2-1 Polarity of Bonds and Molecules

    • 2-1A Bond Dipole Moments

    • Bond polarity ranges: nonpolar covalent, polar covalent, ionic

    • Examples of increasing bond polarity: C–C (nonpolar) < C–N < C–O < C–Cl (more polar) < ionic in methylammonium chloride

    • Bond polarity is quantified by bond dipole moment m

    • Definition: m=qimesdm = q imes d where q is charge separation and d is the distance between charges

    • Units: Debye (D); 1 D = 3.34 × 10^-30 C·m

    • Practical rule of thumb: mext(D)<br>oughly=4.8imesd(extA˚)imesq(extine)m ext{(D)} <br>oughly = 4.8 imes d( ext{Å}) imes q( ext{in }e)

    • Bond dipole moments range from near 0 for symmetric bonds to ~3.6 D for strong polar C≡N

    • The positive end of the dipole points toward the less electronegative atom (partial positive charge)

    • Solved example (2-1): C=O single bond with given μ = 0.86 D and length 1.43 Å

      • Polarization with partial charges: O more electronegative than C

      • Use μ = q d to find charge separation: 0.86=4.8imesd(extA˚)imesq?0.86 = 4.8 imes d( ext{Å}) imes q?

      • Approximated result: q ≈ 0.125 e; i.e., about 1/8 of an electron is effectively separated

      • Conclusion: O bears partial negative charge; C bears partial positive charge (~+0.125 e)

    • Problem 2-1 (b): C=O double bond—μ ≈ 2.4 D, d ≈ 1.23 Å; similarly, compute charge separation and compare resonance contributors (R–C–O− vs R–COH+). (Conceptual: larger μ and shorter d indicate greater charge localization and resonance contribution differences.)

    • 2-1B Molecular Dipole Moments

    • Definition: Molecular dipole moment μ is the vector sum of all bond dipoles in a molecule

    • Symmetry can cancel individual bond dipoles; e.g., CO2 has μ = 0 despite polar C=O bonds due to opposite dipoles canceling

    • Formaldehyde (H2C=O) has μ ≈ 2.3 D; carbon dioxide has μ = 0 D due to symmetry

    • Lone pairs contribute to dipole moments; electron-rich lone pairs increase polarity (e.g., CH2O vs CH2Cl2, etc.)

    • Examples: CH3Cl (μ ≈ 1.9 D), CHCl3 (μ ≈ 1.0–1.9 D depending on geometry), CCl4 (μ = 0 due to tetrahedral symmetry)

    • Problem 2-2: NF3 (m = 0.2 D) vs NH3 (m = 1.5 D) — NF3 has smaller net μ because the strong N–F bonds pull electron density toward F, reducing net dipole; though N–F bond is highly polar, the vector sum yields a small net moment due to geometry and cancellation

    • Problem 2-3 (parts a–e): Use Lewis structures and bond dipoles to evaluate net molecular dipole moments and predict relative polarity

  • 2-2 Intermolecular Forces

    • Noncovalent interactions drive physical properties (melting/boiling points, solubilities)

    • Three major attractive forces:
      1) Dipole–dipole forces (polar molecules)
      2) London dispersion forces (induced, present in all molecules; dominant in nonpolar molecules)
      3) Hydrogen bonding (special dipole–dipole between X–H where X = O, N; water-like interactions)

    • 2-2A Dipole–Dipole Forces

    • Polar molecules align with positive end near negative end of neighbors; overall attraction increases polarity

    • Stronger dipole moments raise heats of vaporization and boiling points for polar compounds

    • Example: CH3Cl engages in dipole–dipole interactions; orientation of dipoles matters for stabilization

    • 2-2B London Dispersion Forces

    • Arise from temporary dipoles induced by momentary electron distribution fluctuations

    • Present in all molecules; become significant in nonpolar species (e.g., CCl4 vs CHCl3): CCl4 can have higher bp despite zero permanent μ due to larger surface area and stronger London forces

    • Boiling point trends in isomeric alkanes: less branching → larger surface area → higher bp (n-pentane > isopentane > neopentane)

    • Geometric surface area correlates with dispersion attraction

    • 2-2C Hydrogen Bonding

    • A hydrogen bond is a strong dipole–dipole attraction, not a true covalent bond

    • Requires H attached to highly electronegative atoms (O–H or N–H; no F–H bonds in typical organic chemistry context here)

    • Polarization: H in O–H or N–H bears a partial positive charge; attracted to lone pairs on electronegative atoms

    • Energy: Hydrogen bonds are strong relative to many other noncovalent forces but far weaker than covalent bonds (~20 kJ/mol vs ~400 kJ/mol for C–H, N–H, O–H covalent bonds)

    • Effect on properties: Ethanol (CH3CH2OH) has BP ≈ 78 °C; Dimethyl ether (CH3OCH3) has BP ≈ -25 °C; ethanol forms H-bonds with water; dimethyl ether does not have O–H, so weaker hydrogen-bonding capacity

    • H-bonding also explains higher boiling points for alcohols than amines with comparable molecular weights; oxygen is more electronegative than nitrogen, leading to stronger O–H interactions

    • Isomer examples: tri-methylamine (no N–H bonds) vs ethylmethylamine vs propylamine; more N–H bonds → higher boiling points due to hydrogen bonding

  • 2-3 Polarity Effects on Solubilities

    • General rule: like dissolves like

    • Four cases:
      1) Polar solute in polar solvent (solvated; dissolution via hydration) — e.g., NaCl in water: ions hydrated by water; large energy release offsets lattice energy; entropy increase favors dissolution
      2) Polar solute in nonpolar solvent — typically does not dissolve well; solvent cannot solvate ions effectively; lattice energy dominates
      3) Nonpolar solute in nonpolar solvent — dissolves due to favorable London dispersion with solvent; entropy increase helps
      4) Nonpolar solute in polar solvent — poor solubility due to disruption of strong solvent–solvent H-bonding networks (hydrophobic effect)

    • Hydration/hydration concept: polar solutes/bonds interact with water via hydration shells; affinity increases dissolution

    • Hydrophobic effect: water forms structured hydrogen-bonding to exclude nonpolar species, leading to poor dissolution and high entropy penalty when forced into solution

    • Application: Biochemistry — vitamins: water-soluble ions; fat-soluble vitamins (A, D) are nonpolar and stored in fat; dose considerations arise from their polarity and solubility

  • 2-4 Arrhenius Acids and Bases

    • Historical definitions:

    • Arrhenius acid: a substance that dissociates in water to give H3O+ (hydronium)

    • Arrhenius base: a substance that dissociates in water to give OH−

    • Brønsted–Lowry refinement (broader applicability):

    • Brønsted–Lowry acid: a species that donates a proton (H+)

    • Brønsted–Lowry base: a species that accepts a proton

    • Includes acids/bases that don’t produce OH− or H3O+ in solution (e.g., NH3 acts as a base toward acids; NH4+ is its conjugate acid)

    • Lewis definitions (broader still):

    • Lewis acid: electron-pair acceptor (electrophile)

    • Lewis base: electron-pair donor (nucleophile)

    • Conjugate acid–base pairs: e.g., NH3/NH4+, H2O/H3O+, CH3OH/CH3O−

    • The important point: many Brønsted–Lowry acids/bases are also Lewis acids/bases

  • 2-5 Brønsted–Lowry Acids and Bases — Conjugates and Examples

    • When acids donate a proton, they form conjugate bases; bases accept a proton and form conjugate acids

    • Notable examples: water as both acid and base; ammonium/ammonia pair; alcohols as both acids and bases under the right conditions

  • 2-6 Strengths of Acids and Bases

    • Strengths quantified by Ka (acid dissociation constant) and Kb (base hydrolysis constant)

    • Strong acids have large Ka (often Ka > 1) and small pKa; weak acids have Ka ≪ 1 and large pKa

    • Common reference values (illustrative):

    • HCl: pKa ≈ -7 (strong acid; fully dissociated in water)

    • H3O+: pKa ≈ -1.7

    • HF: pKa ≈ 3.2 (moderately strong in water context)

    • H2O: pKa ≈ 15.7 (amphiprotic role; water can donate or accept a proton, in self-ionization context it has Kw = 1.0×10^-14)

    • CH3COOH (acetic acid): pKa ≈ 4.74

    • Formic acid (HCOOH): pKa ≈ 3.76

    • HCN: pKa ≈ 9.22

    • Ammonia (NH3): pKa ≈ 9.25 for the conjugate acid NH4+; conjugate acid pKa around 9.25; base strength relates to pKb ≈ 4.75 in water context

    • Methanol (CH3OH): pKa ≈ 15.5–16

    • Key relationship in water: Ka × Kb = Kw; pKa + pKb = 14 (for conjugate pairs in water)

    • Example: Acetic acid + water → CH3COO− + H3O+; compare pKa values to predict equilibrium direction

    • Solvent leveling concept (2-8): In a given solvent, the strongest acids/bases are leveled to the solvent’s own acid/base strength

  • 2-7 Equilibrium Positions of Acid–Base Reactions

    • Brønsted–Lowry reaction transfers a proton from acid (HA) to base (B−): HA + B− → A− + HB

    • Predictions of equilibrium position:

    • The equilibrium favors weaker acid (larger pKa) and weaker base (larger pKb)

    • Weaker acid/conjugate base comparison yields direction

    • If pKa values are known, choose side with higher pKa for acid and higher pKb for base

    • Example: Acetic acid (pKa 4.74) with water (pKa 15.7) — Reaction favors products (weaker acid, aceta te conjugate base) by ~11 pKa units

    • Water’s leveling effect: strong acids in water are leveled to H3O+ (pKa ≈ -1.7) as the strongest acid present; bases are leveled to OH− (pKb ≈ -1.7) in the same sense

    • Practical problem-solving steps (summarized):
      1) Identify acids/bases on both sides
      2) Assign or estimate pKa/pKb values
      3) Use pKa/pKb to compare; the side with the weaker acid or weaker base is favored
      4) If exact pKa/pKb are unavailable, estimate from analogous substances

    • Example problems illustrate using pKa values to determine equilibrium position (e.g., propionic acid vs morpholine; phenol vs aniline)

  • 2-8 Solvent Effects on Acidity and Basicity

    • Water is amphoteric: can act as acid or base; conjugate acid is H3O+, conjugate base is OH−

    • In water, acids stronger than H3O+ are levelled to H3O+; bases stronger than OH− are levelled to OH−

    • pKa scale is solvent-dependent; in water ranges are limited by leveling effects; in other solvents (e.g., diethyl ether, ammonia) different leveling occurs

    • Example solvent comparisons (summary):

    • Water: leveling to H3O+ (pKa -1.7) for acids; to OH− (pKb -1.7) for bases

    • Alcohols: similar leveled ranges; ethanol pKa ≈ 16; ethanol behaves differently as solvent than water

    • Ammonia: liquid ammonia used to run reactions with very strong bases; ammonium/ammonide pair context

    • Ethers: diethyl ether can level acids stronger than pKa ≈ -3.6; ethers do not strongly level bases

    • Alkanes: relatively inert; strong acids/bases in alkanes do not experience leveling

    • Practical takeaway: Solvent choice strongly constrains observed acidity/basicity; Eqn ranges and pKa/pKb values depend on solvent environment

  • 2-9 Effects of Size and Electronegativity on Acidity

    • Acid strength correlates with stability of conjugate base; more stable conjugate base → stronger acid

    • Two primary factors that stabilize conjugate base:

    • Electronegativity: more electronegative atoms stabilize negative charge better

      • Trend: acidity increases from left to right across periodic table

    • Size: larger atoms distribute negative charge better (down the column)

      • Trend: acidity increases down a group

    • Practical takeaway: rightward and downward trends generally increase acid strength

    • Example guidance: fluorine-substituted carboxylic acids are stronger than chlorine-substituted ones due to stronger inductive withdrawal and electronegativity

  • 2-10 Inductive Effects on Acidity

    • Inductive effects: electron-withdrawing atoms/groups stabilize conjugate base through sigma bonds, increasing acidity

    • Dependence on distance (number of bonds between EWG and acidic site): closer substitutions → stronger acid; more substitutions → greater cumulative effect

    • Example illustrations show multi-chloroacetic acid vs mono-chloroacetic acid vs unsubstituted acetic acid

    • Trends demonstrated in pKa charts with halogen substituents (Cl, F) near carboxylate

  • 2-11 Hybridization Effects on Acidity

    • Lone pair stability depends on orbital hybridization where the lone pair resides (on the atom bearing the charge after deprotonation)

    • Higher s-character in the bearing orbital stabilizes the lone pair more (closer to nucleus)

    • Result: acidity increases with increasing s-character in the orbital hosting the lone pair (sp > sp2 > sp3)

    • For nitrogen: sp-hybridized lone-pairs stabilize negative charge more effectively than sp3, increasing acidity of the conjugate base when relevant

    • General rule: acidity order for similar sites: sp3 < sp2 < sp

  • 2-12 Resonance Effects on Acidity and Basicity

    • Resonance delocalization stabilizes conjugate bases, often the dominant effect in increasing acidity

    • Example: acetic acid (pKa ≈ 4.74) vs ethanol (pKa ≈ 15.7) — acetate anion is resonance-stabilized over two oxygens; ethoxide is not

    • Other stabilizing groups: carbonyl (C=O), cyano (C≡N), nitro (NO2) groups strongly stabilize conjugate bases via resonance

    • Amides: resonance delocalization within amide reduces basicity of the amide nitrogen; amides generally far weaker bases than amines (pKb ≈ 14 for acetamide vs ~3.3 for amine)

    • Conceptual point: resonance delocalization contributes significantly to lowering pKa or raising conjugate-base stability; in many cases it dominates over inductive effects

  • 2-13 Lewis Acids and Bases

    • Brønsted–Lowry concepts are a subset of Lewis acid–base definitions

    • Lewis base: electron-pair donor; Lewis acid: electron-pair acceptor

    • Electrophile = Lewis acid; nucleophile = Lewis base

    • Example: NH3 (nucleophile) attacks BF3 (electrophile); curved arrows track electron flow

    • Practical tip: curved arrows depict movement of electron pairs; one arrow per pair of electrons involved

    • Nucleophiles/bases are connected to electron-rich sites; electrophiles/acids are electron-poor sites

  • 2-14 The Curved-Arrow Formalism

    • Formalism for showing electron flow in steps of a reaction

    • Each curved arrow represents the movement of one pair of electrons

    • Arrows originate from lone pairs or from bonds, and point to where electrons will reside after bond formation/bond breaking

    • Color cues are used in the textbook (nucleophiles in blue, electrophiles in green, acidic protons in red) but in notes you can follow standard conventions

    • Practical workflow: use Lewis structures to plan arrows; show bond-forming and bond-breaking events; then analyze resonance processes as needed

  • 2-15 Hydrocarbons (overview of the major classes)

    • Three broad classes for this chapter: hydrocarbons, oxygen-containing compounds, nitrogen-containing compounds

    • Hydrocarbons: C and H only; major subclasses: alkanes (single bonds), alkenes (one C=C), alkynes (one C≡C), arenes (aromatic hydrocarbons, e.g., benzene)

    • General behavior: hydrocarbons are mostly nonpolar or weakly polar; tend to be hydrophobic; typically show London dispersion forces as primary intermolecular force

    • 2-15A Alkanes: single bonds; naming uses -ane; increasing carbons → higher boiling points; branching lowers bp due to reduced surface area

    • 2-15B Alkenes: one C=C; end with -ene; cis/trans isomerism possible; double bonds cannot rotate

    • 2-15D Aromatic Hydrocarbons: arenes (benzene derivatives) with aromatic stabilization; substituents can alter physical properties; Ph denotes phenyl substituent

  • 2-15D Cycloalkanes and Arenes (quick notes)

    • Cycloalkanes: ring forms; cyclopentane, cyclohexane as common examples; naming uses ring parent plus substituents

    • Aromatic hydrocarbons: benzene derivatives; don’t confuse simple cycloalkenes with arenes

    • Toxicology note: methyl/ethyl substitution can reduce overall toxicity relative to benzene due to metabolic processing differences

  • 2-16 Functional Groups with Oxygen

    • 2-16A Alcohols (R–OH)

    • Hydroxy group is strongly polar and can participate in hydrogen bonding; generally more soluble in water if the hydrocarbon portion is not too large

    • Alcohols are miscible with water up to ~4 carbons (ethanol, methanol) and become progressively less water-soluble with more carbons

    • 2-16B Ethers (R–O–R′)

    • Two alkyl groups bonded to oxygen; highly polar but no O–H bond, so no self-hydrogen bonding; ethers do hydrogen-bond with donors (water, alcohols, amines)

    • 2-16C Aldehydes and Ketones (C=O)

    • Carbonyl group is highly polar; can hydrogen-bond with water, alcohols, amines; aldehydes have one alkyl group and one H; ketones have two alkyl groups

    • Solubility in water: up to 4 carbons for miscibility; beyond that, limited solubility

    • 2-16D Carboxylic Acids (R–COOH)

    • Carboxyl group is strongly polar; acidity ~ pKa ≈ 5 in water context (carboxylate stabilization via resonance)

    • Deprotonation yields a resonance-stabilized carboxylate anion (R–COO−)

    • Many common carboxylic acids are miscible with water when small enough; larger carboxylic acids become less soluble

    • 2-16E Carboxylic Acid Derivatives

    • Derivatives include acid chlorides, esters, and amides; all can be hydrolyzed back to carboxylic acids; each has distinct reactivity

    • Examples: acid chloride (R–COCl), ester (R–COOR′), amide (R–CONR′R″)

  • 2-17 Functional Groups with Nitrogen

    • 2-17A Amines

    • Amines are basic; classic structure R–NH2, R2NH, R3N; common alkaloids (e.g., nicotine)

    • Amines typically water-soluble up to about 5 carbons; basicity arises from the lone pair on nitrogen

    • 2-17B Amides

    • Amides are carboxamide derivatives (carboxylic acids converted to CONR′R″); notably stable and common in biology (proteins)

    • Amide nitrogen is less basic than amines due to resonance with the carbonyl; pKb for amide nitrogens is high (weak base)

    • 2-17C Nitriles

    • Nitriles contain the cyano group —C≡N; polar; small nitriles are water-soluble (acetonitrile example)

  • Essential terms (glossary references)

    • acid, base (Arrhenius, Brønsted–Lowry, Lewis definitions)

    • conjugate acid/base, lone pair, electrophile, nucleophile, curved arrows

    • dipole moment, bond dipole moment, molecular dipole moment

    • hydrophilic vs hydrophobic; solubility concepts

    • Ka, pKa, Kb, pKb, Kw, pKw, pH, pH scales

    • resonance, inductive effects, hyperconjugation, hybridization effects

    • common functional groups: alcohol, aldehyde, ketone, carboxylic acid, esters, amides, nitriles, etc.

  • Quick references and problem-solving notes

    • Problem-solving hints: use pKa/pKb to decide equilibria; if pKa difference is large, products are strongly favored

    • For solvent effects: recognize how solvent leveling limits range of measurable pKa/pKb values in that solvent

    • For resonance: resonance stabilization often dominates acidity/basicity; consider stability of conjugate bases and conjugate acids

    • For hybridization: increase s-character to increase acidity of the conjugate base when relevant

    • For Lewis formalism: use curved arrows one per electron-pair transfer; identify electrophiles/nucleophiles clearly

  • Summary of key equations and constants

    • Bond dipole moment (simple): m=qimesdm = q imes d

    • Bond dipole moment units: 1 Debye = 3.34imes10−30extC⋅m3.34 imes 10^{-30} ext{ C·m}

    • Molecular dipole moment: vector sum of bond dipoles (μ is nonzero only when the molecular geometry does not symmetrically cancel all bond dipoles)

    • Dipole approximation: extμ(D)extapproximatelyext=4.8imesqimesd(extA˚)ext{μ (D)} ext{ approximately } ext{=} 4.8 imes q imes d( ext{Å}) when q is in units of e and d is in Å

    • Acid dissociation constant: K<em>a=rac[H</em>3O+][A−][HA]K<em>a = rac{[H</em>3O^+][A^-]}{[HA]}

    • pKa: pK<em>a=−extlog</em>10KapK<em>a = - ext{log}</em>{10} K_a

    • Water ion-product constant: K<em>w=[H</em>3O+][OH−]=1.0imes10−14K<em>w = [H</em>3O^+][OH^-] = 1.0 imes 10^{-14}

    • Brønsted–Lowry conjugate pair concept; pKa + pKb = 14 (in water context) for conjugate acid–base pairs

    • pH: extpH=−extlog<em>10[H</em>3O+]ext{pH} = - ext{log}<em>{10}[H</em>3O^+]

    • Leveling concept: acids stronger than the solvent’s conjugate acid are leveled to that conjugate acid in that solvent; bases stronger than the solvent’s conjugate base are leveled similarly

  • Practical examples and recall cues

    • Alcohols vs ethers: presence of O–H in alcohols enables hydrogen bonding, elevating boiling points and water solubility; ethers lack O–H but can hydrogen-bond with donors, affecting their solubility

    • Carbonyl-containing compounds: carbonyls are highly polar; solubility in water depends on carbon chain length

    • Hydrogen bonding explains large bp differences in isomeric alcohols vs ethers; stronger H-bond donors lead to higher bp

    • The general rule “like dissolves like” is a practical guideline for predicting solubility and miscibility across polar/nonpolar pairs

  • Practical study tips

    • Always check for functional groups and polarity first; that guides solubility and reactivity predictions

    • Use Ka/pKa to rationalize equilibrium directions; if you lack exact values, use analogy to known compounds and then refine with the pKa/pKb relationships

    • When comparing acids, compare conjugate-base stability (resonance, inductive effects, size/electronegativity) rather than just the acid’s structural similarity

    • For Lewis acid–base reactions, practice drawing curved arrows step by step; start by identifying the nucleophile and electrophile, then move electrons accordingly

  • Relevance to real-world chemistry and biology

    • Hydrogen bonding and polarity govern everything from solvent choice in synthesis to protein folding and DNA stability

    • The concept of acidity and basicity influences drug design, enzyme catalysis, and metabolic pathways

    • Inductive and resonance effects explain why small structural changes can drastically alter acidity or basicity, influencing reaction selectivity

  • Quick reference table (condensed)

    • Strong acid in water: Ka large, pKa negative (e.g., HCl, HClO4)

    • Moderate acids: acetic acid (pKa ≈ 4.74), formic acid (pKa ≈ 3.76)

    • Weak acids: acetic acid/ethanol family; pKa around 15–20 range in water context

    • Bases: hydroxide (strong base in water context), amines (moderate bases), amides (weak bases)

    • Common solvents: water, ethanol, diethyl ether, liquid ammonia; pKa/pKb values differ by solvent due to leveling

  • Final takeaway

    • Mastery of Chapter 2 comes from connecting structure to properties: polarity, intermolecular forces, solubility, acidity, basicity, and the different acid–base frameworks (Arrhenius, Brønsted–Lowry, Lewis). The practical tools of this chapter—dipole moments, hydrogen bonding, solubility rules, Ka/pKa, and curved arrows—are the foundation for predicting and rationalizing organic reaction mechanisms and physical properties.