Chapter 2 Acids and Bases - Page-by-Page Notes

Page 1

  • Chapter 2: Acids and Bases. Assigned Reading covers Sections 2.1–2.5 (Slides 1–15). Section 2.11 is OMIT’d.
  • Focus: acids and bases, definitions (Bronsted-Lowry, Lewis, Arrhenius), and the qualitative framework for acid–base strength and equilibrium.

Page 2

  • Chapter Objectives:
    • Identify acids and bases and distinguish between strong and weak acids and bases.
    • Understand how electronegativity, size, hybridization, and resonance affect the stability of bases and how this relates to both base and acid strength.
    • Understand dynamic equilibrium and why the equilibrium is shifted toward the weaker acid–base pair.
    • Evaluate a series of acids or bases and predict their relative acid or base strength.

Page 3

  • Definitions and historical context:
    • Acid (Latin: acidus, sour) — traditionally associated with sour taste.
    • Alkali (base) — from al-qali (ashes of certain plants); bitter taste, slippery feel.
    • Brønsted–Lowry: acid donates a proton (H+); base accepts a proton.
    • Lewis: acid accepts an electron pair; base donates an electron pair.
    • Arrhenius: acids ionize to give H+ in aqueous solution; bases ionize to release OH− in solution (limited use today).

Page 4

  • Acids and Bases (Table 5.2 in the text):
    • Acids provide H+ in aqueous solution. Strong acids completely ionize in water (e.g., HCl, HBr, HI, HNO3, H2SO4, HClO4, Ba(OH)2, LiOH, NaOH, KOH, etc.).
    • Weak acids ionize incompletely; e.g., acetic acid is a weak acid (HA ⇌ H+ + A− with incomplete ionization). The first ionization step may be considered complete in some strong acids (see later sections, e.g., HSO4− in acid-base steps).
    • Bases provide OH− in aqueous solution. Strong bases completely ionize (examples include NaOH, KOH, LiOH, etc.). Weak bases ionize incompletely (e.g., NH3 in water).
    • Note: The image shows acids/bases with approximate symbols (HCl, H2SO4, Ca(OH)2, Sr(OH)2, Ba(OH)2, NaOH, KOH, LiOH, etc.).

Page 5

  • Bases:
    • Bases provide OH− in aqueous solution.
    • Strong bases completely ionize in water: example NaOH(s) → Na+(aq) + OH−(aq).
    • Weak bases incompletely ionize in water (not explicitly listed here but conceptually the reverse of the weak acid equilibrium).

Page 6

  • Recognizing Acids and Bases:
    • Acids have ionizable hydrogen ions (example: CH3CO2H or HC2H3O2).
    • Bases have OH− combined with a metal ion (example: KOH).
    • A representative reaction: Na2CO3(s) + H2O(l) → HCO3−(aq) + 2 Na+(aq) + OH−(aq).
    • This helps identify whether a species is acting as an acid or a base in aqueous solution based on its ability to donate H+ or accept H+.

Page 7

  • Lewis acids and Lewis bases (conceptual definitions):
    • A Lewis base is an electron pair donor.
    • A Lewis acid is an electron pair acceptor.
    • Example interaction: BH3 (a Lewis acid) accepts an electron pair from a Lewis base such as (CH3)3N; the resulting adduct forms via donation of a lone pair to boron.
    • Diagrammatic representation (text): A Lewis base donates an electron pair to a Lewis acid forming a bond.

Page 8

  • More examples of Lewis acids:
    • Common Lewis acids include HCl, HBr, CH3COOH (these are proton donors and can act as Lewis acids).
    • Additional metal-based Lewis acids: BF3, BH3, ZnCl2, AlCl3, FeCl3, TiCl4 (and related compounds).
    • Cations like Li+, Mg2+ can act as Lewis acids in some contexts.

Page 9

  • Lewis bases (examples):
    • Carboxylic acids, amines, sulfides, ethers, alcohols, carboxylate anions, ketones, and other heteroatom-containing species can donate electron pairs.
    • Representative structures reflect donors such as lone pairs on N, O, S, or π-systems that can donate electron density.

Page 10

  • Mechanism: The stronger base wins the tug-of-war for the proton.
    • In the illustrated example, k1 ≫ k−1, so the equilibrium shifts to the right, forming the weaker acid/base pair.
    • The hydroxide ion (OH−) is a much stronger base than the acetate ion (from acetic acid).
    • Net result: formation of the weaker acid/base pair (e.g., OH− + HA ⇌ H2O + A−).
    • This emphasizes the principle that equilibrium favors the formation of the conjugate base of the stronger acid and the conjugate acid of the stronger base, i.e., the weaker conjugate pair is favored.

Page 11

  • Review of Brønsted–Lowry concepts and pH/pKa relationships:
    • Kw: kw=[H+][OH−]=1imes10−14.k_w = [H^+][OH^-] = 1 imes 10^{-14}.
    • pH and pOH relation: pH+pOH=14.pH + pOH = 14.
    • In neutral solution: [H+]=[OH−]=1imes10−7, pH=pOH=7.[H^+] = [OH^-] = 1 imes 10^{-7}, \, pH = pOH = 7.
    • Acid HA in water: HA + H2O ightleftharpoons H3O^+ + A^-. - Equilibrium constant: Ka=[H+][A−][HA].K_a = \frac{[H^+][A^-]}{[HA]}.
    • Base B in water: B + H2O ightleftharpoons BH^+ + OH^-. - Equilibrium constant: K</em>b=[BH+][OH−][B].K</em>b = \frac{[BH^+][OH^-]}{[B]}. - The two constants are related by: K<em>aK</em>b=Kw=1imes10−14.K<em>a K</em>b = K_w = 1 imes 10^{-14}.
    • Henderson–Hasselbalch equations:
    • For acids: pH=pKa+log⁡[A−][HA].pH = pK_a + \log\frac{[A^-]}{[HA]}.
    • For bases: pOH=pKb+log⁡[B][BH+].pOH = pK_b + \log\frac{[B]}{[BH^+]}.
    • Relationship between conjugate pairs: pK<em>a+pK</em>b=14.pK<em>a + pK</em>b = 14. (Note: in the text, pKa and pk1 are related to acid/base conjugates via this Kw relationship.)
    • Practical takeaway: stronger acids have larger Ka and smaller pKa; equilibrium favors weaker acid/base pairs.

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  • pKa scale and classification:
    • Very strong acid: pKa<1.pK_a < 1.
    • Moderately strong acid: pKaextin[1,3].pK_a ext{ in } [1, 3].
    • Weak acid: pKaextin[3,5].pK_a ext{ in } [3, 5].
    • Very weak acid: pKaextin[5,15].pK_a ext{ in } [5, 15].
    • Extremely weak acid: pKa>15.pK_a > 15.
  • pH relation: extpH=− extlog([H+]).ext{pH} = -\ ext{log}([H^+]).
  • The lower the pH, the more acidic the solution; the smaller the pKa, the stronger the acid.

Page 13

  • Approximate pKa ranges for common organic functional groups (Figure 6.3):
    • Carboxylic acids (R–CO2H): ~4–5.
    • Alcohols (R–OH): ~15–16.
    • Protonated alcohols (e.g., ROH2+): highly acidic with pKa near negative values (e.g., ≈ −2.5 for protonated methanol, ≈ −2.4 for protonated ethanol).
    • Protonated amines (RNH3+): ~10.7–11.0 for typical examples.
    • Protonated water (H3O+): pKa ≈ −1.7.
    • Protonated carboxylic acids (e.g., R–CO2H2+ analogs): even more acidic with very negative pKa values (example shown: around −6.1 for certain protonated carboxylic acid species).
    • The figure also indicates a range of pKa values for various organic functional groups to compare acid strength.
  • The general trend: resonance and electronegativity modify acidity; stability of conjugate base is central to acid strength.

Page 14

  • Common organic acids and their pKa values (illustrative examples):
    • Carboxylic acids:
    • Acetic acid, CH3CO2H: pKa ≈ 4.76.
    • Formic acid, HCO2H: pKa ≈ 3.75.
    • Alcohols:
    • Methanol, CH3OH: pKa ≈ 15.5.
    • Ethanol, CH3CH2OH: pKa ≈ 15.9.
    • Protonated alcohols (conjugate acids of alcohols):
    • Protonated methanol (CH3OH2+): pKa ≈ −2.5.
    • Protonated ethanol (CH3CH2OH2+): pKa ≈ −2.4.
    • Amine-related species:
    • Protonated methylamine (CH3NH3+): pKa ≈ 10.7.
    • Ammonia (NH3) as a base has a conjugate acid NH4+ with pKa ≈ 9.4 (approx. values used for comparison in related examples).
    • An example shows methylamine having very different acid/base behavior when protonated vs. deprotonated (pKa values for conjugate acids are around 10.7 for CH3NH3+ and ~11 for related ammonium derivatives).
    • The key takeaway: protonation states and conjugate acids dramatically alter pKa values and acid strength.

Page 15

  • Table 1.8: Approximate pKa values (selected examples represented on the page):
    • H3O+ (water in strongly acidic solution): pKa ≈ −1.7.
    • Protonated water (H3O+) and protonated alcohols show very negative pKa values (e.g., protonated methanol ≈ −2.5; protonated ethanol ≈ −2.4).
    • Carboxylic acids (R–CO2H): pKa around 3.75–4.76 depending on substituents.
    • Neutral alcohols (ROH) have pKa values around 15–16.
    • Protonated amines (RNH3+) show pKa values around 10–11.
    • The table emphasizes the wide span of pKa values across common inorganic and organic species.

Page 16

  • Predicting the direction of acid–base reactions using pKa:
    • Example: Ethanol vs cyanide ion (CN−).
    • Given: pKa(ethanol) ≈ 16; pKa(HCN) ≈ 9.31.
    • Reaction: ext{CH}3 ext{CH}2 ext{OH} + ext{CN}^-
      ightarrow ext{HCN} + ext{CH}3 ext{CH}2 ext{O}^-.
    • Rule: The equilibrium favors the side with the weaker acid (higher pKa). Since ethanol has a higher pKa (weaker acid) than HCN, the left side is favored; the shown right-hand reaction implies a specific context or a particular example trajectory used in the slide (the general rule would predict the left under aqueous conditions). The key concept is using pKa to predict direction.

Page 17

  • Can you neutralize ant bites with NH3 (ammonia)?
    • pKa values involved: NH4+ (conjugate acid of NH3) has pKa ≈ 9.4–9.25; formic acid HCO2H has pKa ≈ 3.75; NH3 acts as a base and can be used to neutralize acids via the formation of ammonium formate: NH3 + HCO2H → NH4+ + HCO2−.
    • The slide shows an acid–base neutralization using formic acid and ammonia, illustrating conjugate acid–base pairs and the direction of proton transfer guided by pKa values.

Page 18

  • Factors determining acid strength (overview):
    1) Electronegativity of the atom bonded to the proton.
    2) Size of the atom bonded to the proton.
    3) Hybridization of the atom bonded to the proton.
    4) Inductive effects of substituents near the proton.
    5) Electron delocalization (resonance) of the conjugate base.
  • Central idea: Relative acid strength correlates with stability of the conjugate base; a more stable conjugate base corresponds to a stronger acid.
  • Consequences: If the conjugate base is stabilized, the corresponding acid is stronger; conversely, stability of base dictates weaker acidity.

Page 19

  • 1) Electronegativity (EN) of the atom bonded to the proton:
    • Increasing EN increases the atom’s ability to attract electrons, stabilizing negative charge on the conjugate base when the proton is removed.
    • Example exercise: Given acids (CH3)3C–H, (CH3)2N–H, CH3O–H, and HF, order from strongest to weakest acid by considering conjugate base stability (the more stable the conjugate base, the stronger the acid).
    • The exercise typically asks to rank acids by the stability of their conjugate bases: A) 1 > 2 > 3 > 4, B) 4 > 3 > 2 > 1, C) 4 > 2 > 3 > 1, D) 1 > 4 > 3 > 2.
  • Concept: Among acids with the same charge and similar size, the one bonded to the more electronegative atom forms a stronger acid due to better stabilization of the conjugate base.

Page 20

  • EN effect on pKa (illustrative examples):
    • CH3OH (methyl alcohol): pKa ≈ 15.5.
    • CH3NH2 (methylamine): pKa for the neutral amine as a base is very high (the conjugate acid CH3NH3+ has pKa ≈ 10.7). In the slide, a value pKa ≈ 40 is shown for CH3NH2 as an acid, illustrating that it is a very weak acid (its conjugate acid is much stronger, with pKa ~ 10.7).
    • Protonated methyl alcohol CH3OH2+ has pKa ≈ −2.5; protonated methylamine CH3NH3+ has pKa ≈ 10.7.
    • Takeaway: When substituents influence EN, they can dramatically shift acidity/basicity by stabilizing or destabilizing the conjugate base.

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  • 2) Size of the atom bonded to the proton:
    • Increasing atomic radius allows the negative charge of the conjugate base to be distributed over a larger volume, stabilizing the anion.
    • Example series: F− < Cl− < Br− < I− (in the same period/group progression, larger halide anions stabilize the conjugate base better, increasing acidity of HX down the group).
    • The table shows an order selection problem: ranking acids H-F, H-Cl, H-Br, H-I.
    • Practical rule: For atoms in the same column, size matters more for conjugate-base stability than electronegativity.

Page 22

  • Potential maps of hydrogen halides illustrate how orbital size correlates with charge distribution and stability of the conjugate base: as orbital volume increases, electron density in the conjugate base becomes more spread out, stabilizing the negative charge.
  • This concept underpins the size effect on acidity for the halogen series.

Page 23

  • Trends in pKa values (summary):
    • When comparing atoms in the same column of the periodic table, size has a greater stabilizing effect on the conjugate base than electronegativity, leading to lower pKa (stronger acid) as size increases.
    • When comparing atoms in the same row, electronegativity is more important than size for determining acidity.
    • Overall: Electronegative atoms contribute to stronger acids, but for same column, larger atoms stabilize the conjugate base more effectively, lowering pKa more than EN would in that column.
  • Concept: “Acid–Base Strength: What Atom is the Charge On?” underscores the importance of conjugate-base stabilization and charge distribution.

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  • Example questions (illustrative):
    • Order compounds by decreasing acidity among options (1, 2, 3) with a given set of substituents (O, F, Cl, etc.).
    • Identify which hydrogens in a molecule are most acidic based on potential resonance/stabilization of the conjugate base.
  • Core idea: Use conjugate-base stability (including inductive and resonance effects) to rank acidity and to identify the most acidic H in a given structure.

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  • 3) Hybridization of the atom bonded to the proton:
    • Acid strength correlates with hybridization: sp > sp2 > sp3.
    • This trend reflects the percentage of s-character in the C–H (or X–H) bond and the resulting ability to stabilize negative charge on the conjugate base after deprotonation.
    • In terms of s-character: sp has 50% s-character, sp2 has 33%, sp3 has 25%.

Page 26

  • 4) Inductive Effects (through-bond electron withdrawal):
    • Inductive effect: shifting of electrons in a sigma bond in response to nearby electronegative atoms.
    • Electron withdrawal inductively stabilizes the conjugate base, increasing acidity: HA \/ A− + H+.
    • Inductive effects propagate through sigma bonds; electron-withdrawing substituents near the proton make the carbon more positive and stabilize the conjugate base.
    • Location of substituents matters: closer to the proton yields a stronger inductive effect.

Page 27

  • Details on Inductive Effects:
    • Inductive electron withdrawal increases acidity; greater EN and closer proximity yield stronger acidity.
    • Example statement: Inductive effects are transmitted along sigma bonds; adding multiple electron-withdrawing groups increases acidity by decreasing the basicity of the conjugate base.
    • Relative positioning of substituents affects acid strength (proximal substituents have larger effects).

Page 28

  • Quantitative illustration of inductive effects on pKa:
    • Values listed: pKa around 4.75, 2.87, 1.25, 0.70 for successive examples of inductive withdrawal.
    • In carboxylic acids, proximity and number of electron-withdrawing substituents near the CO2H group strongly influence acidity.
  • Examples of substituted carboxylic acids showing pKa differences due to inductive effects:
    • Carboxylic acids with various substituents can give pKa values such as 4.83, 4.52, 4.05, 2.86 (illustrating that electron-withdrawing groups near CO2H increase acidity).
  • Summary: Inductive effects can dramatically alter pKa by stabilizing the conjugate base through sigma-bond electron withdrawal.

Page 29

  • Solvation effects in water:
    • The solvent stabilizes the negatively charged conjugate base through solvation.
    • If the conjugate base is poorly solvated (unfavorable solvent interactions), the equilibrium shifts toward the reactants, making the acid weaker.
    • Carboxylic acids illustrate this: more sterically hindered acids are less solvated and thus less acidic than their less hindered counterparts.
    • Example pKa patterns for carboxylic acids and comparable alcohols show that solvation can shift acidity meaningfully.
  • General trend: Solvation (hydrogen-bonding and dielectric effects) stabilizes anions and shifts equilibria toward the side with the more stabilized conjugate base.

Page 30

  • Acid–Base Strength of alkyl-substituted acids and solvent effects:
    • Solvation differences influence acidity: stabilizing the conjugate base via solvation shifts equilibrium to the right for more solvated conjugate bases.
    • Example families: formic acid (pKa ≈ 3.75) vs 1,1-dimethylpropanoic acid (pKa ≈ 5.03) vs other substituted carboxylic acids (examples shown include pKa ≈ 4.76 for acetic acid, ≈ 4.85 for isobutyric-like acids).
    • Solvent effects (e.g., water as solvent) influence acidity via differential stabilization of A−.
    • Hydrophilic–hydrophobic interactions can be unfavorable for solvation, altering observed acidity.

Page 31

  • Practice concept questions (illustrative):
    • Order acids from most acidic to least acidic based on conjugate-base stability and solvation considerations (options A–E given).
    • prompted to draw conjugate bases and analyze resonance/inductive effects before ranking.
    • The key skill: apply the stability of the conjugate base (delocalization, inductive effects, and solvation) to rank acidity.

Page 32

  • Electron delocalization and acidity (delocalization resonance effects):
    • Compare CH3–CH2–H vs CH3–CH2–O–H vs CH3–C(=O)–H (carboxylic acid) or related structures.
    • Acid strengths: acetic acid is stronger than ethanol due to resonance stabilization of the acetate anion.
    • Resonance structures stabilize negative charge on the conjugate base by delocalizing charge over oxygen atoms in carboxylates, versus less stabilization for non-resonance-containing species.
    • The example shows that CH3CO2H has pKa ≈ 4.74, while ethanol and related alcohols have much higher pKa values (less acidic).
    • Delocalization (resonance) is a major driver of conjugate-base stability and thus acid strength.

Page 33

  • Visual comparison of resonance and acidity:
    • CH3–CH2–H (alkane) has minimal resonance stabilization of the conjugate base; thus weaker acidity.
    • CH3–CH2–O–H (alcohol) has some stabilization but less than carboxylates.
    • CH3CO2H (acetic acid) demonstrates strong conjugate-base stabilization via resonance, hence stronger acidity.
    • The slide shows pKa values: CH3CH3 (likely a misprint here; intended to compare CH3CH2– with resonance vs non-resonance cases) and pKa values such as 15.9 for ethanol, 4.76 for acetic acid, and 51 for a resonance-limited case (illustrative).
  • Core concept: Electron delocalization stabilizes conjugate bases and increases acid strength.

Page 34

  • Delocalization of charge and resonance in acidity:
    • No resonance vs resonance cases:
    • No resonance: higher energy conjugate base, weaker acid.
    • Resonance stabilization (e.g., in carboxylate anions) dramatically stabilizes the conjugate base and increases acidity.
    • Examples:
    • CH3CO2H (acetic acid): pKa ≈ 4.74.
    • CH3O− + H+ ⇌ CH3OH pKa values illustrate resonance stabilization in different contexts.
    • The overall message: Delocalization of charge (resonance) lowers the energy of the conjugate base, thereby increasing acid strength.

Page 35

  • Connecting delocalization, substituent effects, and pKa in drug design and reaction mechanisms:
    • Which phenol is the stronger acid? The presence and position of electron-withdrawing substituents (e.g., NO2) on the aromatic ring stabilizes the phenoxide anion via resonance, increasing acidity.
    • Visualization: a phenol with two nitro groups (NO2) adjacent to the OH group is typically more acidic due to resonance stabilization of the resulting phenoxide.

Page 36

  • Example structure: a phenol with NO2 substituents (depicted) to illustrate the effect of electron-withdrawing groups on acidity.

Page 37

  • Delocalization of charge and acidity: a practice problem
    • Which hydrogen is most acidic (1 or 2)?
    • Approach: Draw possible conjugate bases; draw resonance structures for each conjugate base; resonance stabilizes the anion by delocalizing the charge across atoms, making that conjugate base more stable and the corresponding proton more acidic.

Page 38

  • Repeated problem: Which hydrogens (1 or 2) would be most acidic? Answer depends on the ability to delocalize negative charge via resonance in the resulting conjugate base.
  • Emphasizes resonance stabilization as a determinant of acidity.

Page 39

  • Summary: Factors affecting acid strength
    • 1) Size
    • 2) Electronegativity
    • 3) Hybridization
    • 4) Inductive effect
    • 5) Electron delocalization (resonance)

Page 40

  • Key Concepts (recap):

    • Acid vs. Brønsted–Lowry acid: a species that donates a proton. Brønsted–Lowry base: a species that accepts a proton.
    • Lewis acid: accepts an electron pair; Lewis base: donates an electron pair.
    • Ka and pKa quantify acid strength; stronger acids have larger Ka and smaller pKa.
    • In acid–base reactions, equilibria favor formation of the weaker acid and weaker base (i.e., the conjugate pair with higher stability).
    • Relationship between pKa and pH: pH is the negative log of the hydrogen ion concentration; the pKa is the pH at which the acid is half dissociated.
    • The stability of the conjugate base dictates acid strength; similar-sized atoms: the more electronegative the atom attached to H, the stronger the acid. When atoms differ in size, the larger atom attached to H often leads to a stronger acid due to better charge dispersal.
    • Hybridization: sp > sp2 > sp3 in determining acidity, with greater s-character stabilizing the conjugate base.
    • Inductive effects: electron withdrawal through sigma bonds increases acidity, especially when the withdrawing groups are close to the proton.
    • Electron delocalization (resonance) increases conjugate-base stability and acid strength.
    • pKa–pH relationship: in acidic solution, the compound tends to stay in its acidic form; in basic solution, it tends to be in its basic form.

  • Connections to foundational principles:

    • The concepts tie directly to thermodynamics of acid–base equilibria (Le Châtelier’s principle via stabilization of conjugate bases).
    • Practical relevance includes predicting reaction directions, stability of intermediates, and design considerations in organic synthesis and medicinal chemistry.