Acidity, Basicity, and pKa (UG/PG)

Chapter: Acidity, Basicity, and pKa

1. Introduction to Acids and Bases

  • Arrhenius Definition:
    • Acid: A substance that produces hydrogen ions (H+) in aqueous solution.
    • Base: A substance that produces hydroxide ions (OH−) in aqueous solution.
    • Limitation: Restricted to aqueous solutions.
  • Brønsted-Lowry Definition:
    • Acid: A proton (H+) donor.
    • Base: A proton (H+) acceptor.
    • Conjugate Acid-Base Pairs: When a Brønsted-Lowry acid donates a proton, it forms its conjugate base. When a Brønsted-Lowry base accepts a proton, it forms its conjugate acid.
      • Example: HA+B⇌A−+BH+
        • HA is the acid, A− is its conjugate base.
        • B is the base, BH+ is its conjugate acid.
    • Key Concept: Acid-base reactions involve the transfer of a proton.
  • Lewis Definition:
    • Acid: An electron pair acceptor (electrophile).
    • Base: An electron pair donor (nucleophile).
    • Lewis Adduct: The product formed from a Lewis acid-base reaction.
    • Broadest Definition: Encompasses Brønsted-Lowry acids/bases and reactions not involving protons (e.g., BF3​ reacting with NH3​).

2. Acid Strength and pKa

  • Acid Strength: A measure of how readily an acid donates a proton. Stronger acids donate protons more readily.
  • Equilibrium Constant (Ka​): For a generic acid HA in water:
    HA(aq)+H2​O(l)⇌A−(aq)+H3​O+(aq)Ka​=[HA][A−][H3​O+]​
    • A larger Ka​ value indicates a stronger acid.
  • pKa: The negative logarithm (base 10) of the Ka​ value:
    pKa=−log10​Ka​
    • A smaller pKa value (or more negative) indicates a stronger acid.
    • A larger pKa value indicates a weaker acid.
    • Rule of Thumb: A strong acid has a weak conjugate base, and a weak acid has a strong conjugate base.
  • Predicting Reaction Direction: An acid-base reaction will favor the formation of the weaker acid and weaker base. The equilibrium will lie on the side of the higher pKa value.
    • Example: HA (pKa 5) + B− (pKa of HB is 10) ⇌A−+HB. The equilibrium favors HB and A− because HB is the weaker acid.

3. Factors Affecting Acidity (and Conjugate Base Stability)

The strength of an acid is primarily determined by the stability of its conjugate base (A−). A more stable conjugate base corresponds to a stronger acid.

  • Electronegativity (Across a Row):
    • Acidity increases across a row in the periodic table (e.g., CH4​<NH3​<H2​O<HF).
    • The more electronegative the atom bearing the negative charge in the conjugate base, the better it can stabilize that charge, thus increasing acidity. (e.g., F− is more stable than O2− or N3−).
  • Atomic Size/Polarizability (Down a Group):
    • Acidity increases down a group in the periodic table (e.g., HF<HCl<HBr<HI).
    • Even though electronegativity decreases down a group, the increasing size (and thus polarizability) of the atom bearing the negative charge in the conjugate base allows for better delocalization of the charge over a larger volume, stabilizing it.
  • Hybridization:
    • Acidity of C-H bonds: sp (≡C-H) > sp2 (=C-H) > sp3 (-C-H).
    • The higher the ‘s’ character in the hybridized orbital, the closer the lone pair in the conjugate base is to the nucleus. This leads to better stabilization of the negative charge and increased acidity.
      • Example: Terminal alkynes are weakly acidic (pKa ≈25) due to the sp hybridized carbon, allowing them to be deprotonated by strong bases like Grignard or organolithium reagents.
  • Inductive Effects:
    • Electron-Withdrawing Groups (EWGs): Stabilize the negative charge on the conjugate base by pulling electron density away through sigma bonds. This increases acidity. The closer the EWG to the acidic proton, the stronger the effect.
      • Example: Chloroacetic acid is stronger than acetic acid (CH3​COOH vs. ClCH2​COOH).
    • Electron-Donating Groups (EDGs): Destabilize the negative charge on the conjugate base by pushing electron density towards it. This decreases acidity.
      • Example: Tert-butanol is a weaker acid than methanol.
  • Resonance Effects:
    • If the negative charge in the conjugate base can be delocalized over multiple atoms through resonance, the conjugate base is more stable, and the parent acid is stronger.
    • Example: Carboxylic acids are much stronger acids than alcohols because the carboxylate anion’s negative charge is delocalized over two oxygen atoms. Phenols are more acidic than cyclohexanol because the phenoxide anion’s negative charge is delocalized into the aromatic ring.
    • Aromaticity: If the formation of the conjugate base leads to an aromatic system (e.g., cyclopentadiene), the acid is significantly more acidic.
  • Solvation Effects:
    • The ability of the solvent to stabilize the charged conjugate base through solvation (interaction with solvent molecules) can significantly impact acidity. Protic solvents (like water, alcohols) can stabilize anions through hydrogen bonding.
    • A more stabilized conjugate base by solvation implies a stronger acid.

4. Factors Affecting Basicity

The strength of a base is primarily determined by its ability to accept a proton (i.e., the availability of its lone pair of electrons to form a new bond with a proton). A stronger base has a higher affinity for protons.

  • Electronegativity (Across a Row):
    • Basicity decreases across a row in the periodic table (e.g., CH3−​>NH2−​>OH−>F−).
    • The more electronegative the atom bearing the lone pair, the less willing it is to share that lone pair, thus decreasing basicity.
  • Atomic Size/Polarizability (Down a Group):
    • Basicity decreases down a group (e.g., F−>Cl−>Br−>I−).
    • While larger atoms are more polarizable, which helps stabilize a negative charge (making the conjugate acid stronger), the lone pair on a larger atom is less concentrated and thus less available to accept a proton effectively.
  • Hybridization:
    • Basicity of nitrogen atoms: sp3 (amines) > sp2 (imines/pyridines) > sp (nitriles).
    • Higher ‘s’ character means the lone pair is held closer to the nucleus, making it less available for protonation, decreasing basicity.
      • Example: Pyrrole (nitrogen lone pair part of aromatic system) is a very weak base compared to pyridine (nitrogen lone pair in sp2 orbital, not part of aromaticity).
  • Resonance Effects:
    • If the lone pair of electrons on a potential basic atom is delocalized through resonance, it is less available to accept a proton, thus decreasing basicity.
    • Example: Anilines are weaker bases than aliphatic amines because the nitrogen lone pair is delocalized into the aromatic ring. Amides are very weak bases because the nitrogen lone pair is delocalized onto the carbonyl oxygen.
  • Inductive Effects:
    • Electron-Donating Groups (EDGs): Increase electron density on the basic atom, making the lone pair more available for protonation, increasing basicity.
      • Example: Methylamine (CH3​NH2​) is more basic than ammonia (NH3​). Alkyl groups are typically EDGs.
    • Electron-Withdrawing Groups (EWGs): Decrease electron density on the basic atom, making the lone pair less available for protonation, decreasing basicity.
      • Example: Trifluoroethylamine is a weaker base than ethylamine.
  • Steric Effects:
    • Very bulky groups around a basic site can hinder the approach of a proton, thereby decreasing basicity (steric hindrance to protonation). This is often observed in the gas phase, but less so in solution due to solvation effects.

5. Practical Considerations

  • Strong Acids: Acids with pKa values lower than that of H3​O+ (pKa = -1.74) are considered strong acids in water. They completely dissociate.
  • Strong Bases: Bases with conjugate acids having pKa values higher than that of H2​O (pKa = 15.7) are considered strong bases in water. They will completely deprotonate water.
  • Leveling Effect: In a given solvent, all acids stronger than the solvent’s conjugate acid will appear to have the same strength (they all fully protonate the solvent). Similarly, all bases stronger than the solvent’s conjugate base will appear to have the same strength (they all fully deprotonate the solvent). This limits the range of pKa that can be measured in a particular solvent.
  • Solvents: The choice of solvent can drastically alter observed pKa values due to differences in solvation.

6. Common Functional Groups and their Approximate pKa Values

  • Strong Acids (pKa < 0): HCl (-7), H2​SO4​ (-3), HNO3​ (-1.4)
  • Carboxylic Acids: pKa ≈3−5 (e.g., Acetic acid, 4.76)
  • Phenols: pKa ≈10 (e.g., Phenol, 9.95)
  • Alcohols: pKa ≈16−18 (e.g., Ethanol, 16)
  • Water: pKa = 15.7
  • Terminal Alkynes: pKa ≈25
  • Amines (as conjugate acids RNH3+​): pKa ≈9−11 (e.g., Methylammonium ion, 10.6)
  • Ammonia: pKa ≈36 (as NH3​ acting as an acid, not as conjugate acid)
  • Alkanes: pKa > 50

Multiple Choice Questions (MCQ) on Acidity, Basicity, and pKa

Instructions: Choose the best answer for each question.

1. According to the Arrhenius definition, what does an acid produce in aqueous solution? a) Hydroxide ions (OH−) b) Hydrogen ions (H+) c) A precipitate d) A salt

2. Which definition describes a base as a proton (H+) acceptor? a) Arrhenius b) Brønsted-Lowry c) Lewis d) None of the above

3. In the reaction HA+B⇌A−+BH+, what is A−? a) The acid b) The base c) The conjugate acid of HA d) The conjugate base of HA

4. A Lewis acid is defined as a(n): a) Proton donor b) Proton acceptor c) Electron pair donor d) Electron pair acceptor

5. Which of the following pKa values represents the strongest acid? a) 5.2 b) 1.5 c) -3.0 d) 10.1

6. If a reaction has an acid with a pKa of 4 and a conjugate acid with a pKa of 9, which side will the equilibrium favor? a) Reactants b) Products c) Neither, it will be at equilibrium d) Cannot be determined without more information

7. Which factor explains why HF is a weaker acid than HCl? a) Electronegativity b) Atomic size/Polarizability c) Hybridization d) Resonance

8. Based on electronegativity, arrange the following in increasing order of acidity: CH4​, NH3​, H2​O. a) CH4​<NH3​<H2​O b) H2​O<NH3​<CH4​ c) NH3​<CH4​<H2​O d) CH4​<H2​O<NH3​

9. Why are terminal alkynes significantly more acidic than alkanes? a) Due to resonance stabilization of the conjugate base. b) Due to the sp hybridization of the carbon bearing the acidic hydrogen. c) Due to inductive electron-donating effects. d) Due to strong hydrogen bonding.

10. Which of the following groups, when attached to a molecule, would increase its acidity via an inductive effect? a) −CH3​ (methyl) b) −F (fluoro) c) −OH (hydroxyl) d) −NH2​ (amino)

11. Why is acetic acid (CH3​COOH) a stronger acid than ethanol (CH3​CH2​OH)? a) Acetic acid has more hydrogen atoms. b) The conjugate base of acetic acid is stabilized by resonance. c) Ethanol has a higher boiling point. d) Acetic acid is a heavier molecule.

12. Which of the following will be the strongest base? a) Cl− b) Br− c) I− d) F−

13. Why is ammonia (NH3​) a stronger base than pyridine? a) The lone pair in pyridine is part of an aromatic system. b) The nitrogen in pyridine is sp2 hybridized, making its lone pair less available. c) Ammonia has more hydrogen atoms. d) Pyridine is a bulkier molecule.

14. Which of the following would decrease the basicity of an amine via an inductive effect? a) Alkyl groups attached to nitrogen. b) Electron-donating groups near the nitrogen. c) Electron-withdrawing groups near the nitrogen. d) Aromatic rings attached to nitrogen.

15. What is the approximate pKa of a typical carboxylic acid? a) -5 b) 4-5 c) 10 d) 16-18

16. What is the approximate pKa of water? a) 0 b) 7 c) 15.7 d) 36

17. Which of the following is considered a strong acid in water (i.e., fully dissociates)? a) Acetic acid (pKa 4.76) b) Phenol (pKa 9.95) c) HCl (pKa -7) d) Water (pKa 15.7)

18. What is the leveling effect of a solvent? a) All acids appear to have the same strength in that solvent if they are stronger than the solvent’s conjugate acid. b) All bases appear to have the same strength if they are weaker than the solvent. c) It makes all acids appear weaker than they actually are. d) It only applies to non-polar solvents.

19. A compound with a pKa of 25 is best classified as a: a) Strong acid. b) Weak acid. c) Strong base. d) Neutral compound.

20. Which of the following is a Lewis acid but not typically a Brønsted-Lowry acid? a) HCl b) H2​SO4​ c) BF3​ d) H3​O+

21. What happens to the acidity of an alcohol if an electron-withdrawing group is placed near the hydroxyl group? a) Acidity decreases. b) Acidity increases. c) Acidity remains unchanged. d) It becomes basic.

22. Which of the following is the strongest base among the conjugate bases derived from the acids listed: HF, H2​O, NH3​, CH4​? a) F− b) OH− c) NH2−​ d) CH3−​

23. The acidity of carboxylic acids compared to alcohols is primarily due to: a) Steric effects. b) Inductive effects. c) Resonance stabilization of the carboxylate anion. d) Hydrogen bonding in alcohols.

24. Which type of compound is generally the most basic among the following? a) Amides b) Anilines c) Aliphatic amines d) Pyrroles

25. A substance that accepts a proton is a Brønsted-Lowry: a) Acid b) Base c) Conjugate acid d) Conjugate base

26. What is the relationship between the strength of an acid and the strength of its conjugate base? a) Strong acid, strong conjugate base. b) Weak acid, weak conjugate base. c) Strong acid, weak conjugate base. d) No direct relationship.

27. Which of the following factors contributes to the increased acidity of phenols compared to cyclohexanol? a) Inductive effect of the phenyl group. b) Resonance stabilization of the phenoxide anion into the aromatic ring. c) Angle strain in cyclohexanol. d) Hydrogen bonding in phenol.

28. If an acid has a very large Ka​ value, its pKa will be: a) A very large positive number. b) Close to zero. c) A very small or negative number. d) Cannot be determined.

29. The pKa of a proton is 10. The pKa of another proton is 5. If we react the conjugate base of the pKa 10 acid with the acid of the pKa 5, which species will be favored at equilibrium? a) The pKa 10 acid and its conjugate base. b) The pKa 5 acid and its conjugate base. c) The pKa 10 acid and the pKa 5 acid. d) The pKa 5 conjugate base and the pKa 10 conjugate acid.

30. Which solvent would be best for observing the full difference in acidity between two very strong acids, like HCl and HBr? a) Water b) Ethanol c) Glacial acetic acid d) Diethyl ether

31. The basicity of an amine is decreased if the nitrogen’s lone pair is: a) Delocalized by resonance. b) Part of a saturated ring. c) Exposed to steric hindrance. d) Participating in hydrogen bonding with solvent.

32. What type of effect is responsible for the observation that chloroacetic acid is a stronger acid than acetic acid? a) Resonance effect. b) Inductive effect. c) Hyperconjugation. d) Steric effect.

33. Rank the following C-H bonds in increasing order of acidity: alkane, alkene, alkyne. a) alkane<alkene<alkyne b) alkyne<alkene<alkane c) alkene<alkane<alkyne d) alkane<alkyne<alkene

34. The stability of a conjugate base is directly related to the strength of its corresponding acid. A more stable conjugate base means: a) A weaker acid. b) A stronger acid. c) An unreactive acid. d) A neutral acid.

35. Which definition of acids and bases is most useful in organic chemistry because it is the broadest and applies to reactions not involving protons? a) Arrhenius b) Brønsted-Lowry c) Lewis d) All are equally useful

36. If a functional group has an approximate pKa of 10, it is likely a(n): a) Carboxylic acid. b) Alcohol. c) Phenol. d) Alkane.

37. How does increased ‘s’ character in a hybrid orbital affect the basicity of an atom? a) Increases basicity. b) Decreases basicity. c) Has no effect. d) Only affects acidity, not basicity.

38. Which of the following would be the strongest base? a) Methylamine (CH3​NH2​) b) Aniline (C6​H5​NH2​) c) Acetamide (CH3​CONH2​) d) Pyrrole (C4​H4​NH)

39. What is the role of the solvent in stabilizing a conjugate base through hydrogen bonding? a) It makes the acid weaker. b) It makes the conjugate base less reactive. c) It helps delocalize the negative charge on the conjugate base. d) It increases the basicity of the conjugate base.

40. If a base is stronger than OH−, it will typically react completely with water to form OH−. This is an example of the: a) Equilibrium constant. b) Leveling effect. c) Inductive effect. d) Resonance effect.

Answer Key with Explanations

  1. b) Hydrogen ions (H+).
    • Explanation: The Arrhenius definition states that an acid is a substance that dissociates in water to produce hydrogen ions (H+).
  2. b) Brønsted-Lowry.
    • Explanation: The Brønsted-Lowry definition defines an acid as a proton donor and a base as a proton acceptor.
  3. d) The conjugate base of HA.
    • Explanation: When an acid (HA) donates a proton, the species remaining (A−) is its conjugate base.
  4. d) Electron pair acceptor.
    • Explanation: A Lewis acid is defined as an electron pair acceptor, while a Lewis base is an electron pair donor.
  5. c) -3.0.
    • Explanation: A smaller (or more negative) pKa value indicates a stronger acid. Out of the given options, -3.0 is the smallest.
  6. b) Products.
    • Explanation: Acid-base reactions favor the formation of the weaker acid and weaker base. The equilibrium will lie on the side of the higher pKa value (in this case, pKa 9 for the conjugate acid).
  7. b) Atomic size/Polarizability.
    • Explanation: While F is more electronegative than Cl, the larger size of the chlorine atom allows it to better delocalize and stabilize the negative charge in its conjugate base (Cl−) compared to F−. This effect dominates down a group.
  8. a) CH4​.
    • Explanation: Acidity increases across a row in the periodic table due to increasing electronegativity. The more electronegative the atom bearing the negative charge in the conjugate base, the more stable it is.
  9. b) Due to the sp hybridization of the carbon bearing the acidic hydrogen.
    • Explanation: The sp hybridized carbon in a terminal alkyne has 50% ‘s’ character. This higher ‘s’ character means the electrons in the C-H bond are held closer to the nucleus, making the hydrogen more acidic and the resulting carbanion more stable.
  10. b) −F (fluoro).
    • Explanation: Electron-withdrawing groups (EWGs) like fluorine stabilize the negative charge of the conjugate base through the inductive effect (pulling electron density through sigma bonds), thus increasing acidity.
  11. b) The conjugate base of acetic acid is stabilized by resonance.
    • Explanation: The carboxylate anion (CH3​COO−) formed from acetic acid can delocalize its negative charge over two oxygen atoms through resonance, making it much more stable than the alkoxide ion from ethanol.
  12. d) F−.
    • Explanation: Basicity decreases down a group. Since HF is the weakest acid among the HX acids (due to poor charge delocalization in F−), its conjugate base, F−, will be the strongest base.
  13. b) The nitrogen in pyridine is sp2 hybridized, making its lone pair less available.
    • Explanation: The lone pair on the nitrogen in pyridine is in an sp2 orbital, which has more ‘s’ character (and is thus closer to the nucleus and less available for protonation) than the sp3 hybridized lone pair in ammonia. Also, the lone pair in pyridine is not part of the aromatic system, unlike pyrrole.
  14. c) Electron-withdrawing groups near the nitrogen.
    • Explanation: Electron-withdrawing groups decrease the electron density on the basic atom (nitrogen), making its lone pair less available for protonation and thus decreasing basicity.
  15. b) 4-5.
    • Explanation: Carboxylic acids typically have pKa values ranging from approximately 3 to 5, making them moderately acidic.
  16. c) 15.7.
    • Explanation: The pKa of water is approximately 15.7. This is a crucial reference point for acid-base strength in aqueous solutions.
  17. c) HCl (pKa -7).
    • Explanation: Strong acids are those with pKa values lower than the pKa of H3​O+ (-1.74). HCl falls into this category, meaning it fully dissociates in water.
  18. a) All acids appear to have the same strength in that solvent if they are stronger than the solvent’s conjugate acid.
    • Explanation: The leveling effect describes how the strength of an acid (or base) is limited by the solvent’s ability to be protonated (or deprotonated). In water, any acid stronger than H3​O+ will simply fully protonate water to form H3​O+, appearing to have the same strength.
  19. b) Weak acid.
    • Explanation: A pKa of 25 is a relatively high value, indicating a very weak acid (and consequently, a very strong conjugate base). For context, water has a pKa of 15.7.
  20. c) BF3​.
    • Explanation: BF3​ is an electron-deficient compound (lacks an octet) and can accept an electron pair, making it a Lewis acid. However, it does not have a proton to donate, so it is not a Brønsted-Lowry acid.
  21. b) Acidity increases.
    • Explanation: Electron-withdrawing groups stabilize the negative charge on the conjugate base (alkoxide ion) by dispersing electron density, making the alcohol a stronger acid.
  22. d) CH3−​.
    • Explanation: The weakest acid among the given options is CH4​. Therefore, its conjugate base, CH3−​, will be the strongest base. This follows the inverse relationship between acid and conjugate base strength.
  23. c) Resonance stabilization of the carboxylate anion.
    • Explanation: The key difference is that the negative charge on the carboxylate ion can be delocalized over two electronegative oxygen atoms, making it far more stable than the localized negative charge on the alkoxide ion from an alcohol.
  24. c) Aliphatic amines.
    • Explanation: Aliphatic amines typically have the nitrogen lone pair readily available for protonation. Amides and anilines have their lone pairs delocalized by resonance, reducing basicity. Pyrrole’s nitrogen lone pair is part of its aromatic system, making it very weakly basic.
  25. b) Base.
    • Explanation: By definition, a Brønsted-Lowry base is a proton acceptor.
  26. c) Strong acid, weak conjugate base.
    • Explanation: There is an inverse relationship: the stronger an acid, the weaker its conjugate base, and vice-versa.
  27. b) Resonance stabilization of the phenoxide anion into the aromatic ring.
    • Explanation: When phenol loses a proton, the resulting phenoxide anion can delocalize its negative charge into the benzene ring through resonance, stabilizing the anion and making phenol more acidic than a simple alcohol like cyclohexanol.
  28. c) A very small or negative number.
    • Explanation: pKa = -log(Ka​). If Ka​ is very large (e.g., 107), then pKa will be a large negative number (e.g., -7).
  29. b) The pKa 5 acid and its conjugate base.
    • Explanation: The reaction proceeds from the stronger acid/stronger base pair to the weaker acid/weaker base pair. The acid with pKa 5 is stronger than the acid with pKa 10. Therefore, the equilibrium will favor the formation of the weaker acid (pKa 10 acid) and its conjugate base. So, the pKa 5 acid will donate its proton to the conjugate base of the pKa 10 acid.
  30. c) Glacial acetic acid.
    • Explanation: To observe the difference in strength between two very strong acids, you need a solvent that is a weaker base than water (and therefore, its conjugate acid has a lower pKa than H3​O+). Glacial acetic acid is less basic than water and can differentiate between strong acids (it’s a “differentiating solvent”).
  31. a) Delocalized by resonance.
    • Explanation: If the nitrogen’s lone pair is involved in resonance, it is less available to accept a proton, thus decreasing the basicity of the amine (e.g., anilines, amides).
  32. b) Inductive effect.
    • Explanation: The highly electronegative chlorine atom in chloroacetic acid pulls electron density away from the carboxylate oxygen through sigma bonds (inductive withdrawal), stabilizing the negative charge and making it a stronger acid than acetic acid.
  33. a) alkane.
    • Explanation: Acidity of C-H bonds increases with increasing s-character of the carbon orbital: sp3 (alkane, 25% s) < sp2 (alkene, 33% s) < sp (alkyne, 50% s).
  34. b) A stronger acid.
    • Explanation: The more stable the conjugate base, the more readily the acid will donate its proton, making it a stronger acid.
  35. c) Lewis.
    • Explanation: The Lewis definition (electron pair donor/acceptor) is the broadest definition and is particularly useful in organic chemistry because it encompasses a wide range of reactions, including those that do not involve proton transfer.
  36. c) Phenol.
    • Explanation: Phenols typically have pKa values around 10. Carboxylic acids are much lower (4-5), alcohols are higher (16-18), and alkanes are much higher (>50).
  37. b) Decreases basicity.
    • Explanation: Higher ‘s’ character means the lone pair electrons are held more closely to the nucleus (due to the s orbital’s spherical shape and lower energy). This makes them less available for donation to a proton, thus decreasing basicity.
  38. a) Methylamine (CH3​NH2​).
    • Explanation: Methylamine is an aliphatic amine. Its basicity is enhanced by the electron-donating inductive effect of the methyl group. Aniline and acetamide have resonance-delocalized lone pairs, and pyrrole’s lone pair is part of an aromatic system, all making them significantly weaker bases.
  39. c) It helps delocalize the negative charge on the conjugate base.
    • Explanation: Protic solvents can form hydrogen bonds with the negatively charged conjugate base, which helps to stabilize and disperse the charge, thus increasing the acidity of the original acid.
  40. b) Leveling effect.
    • Explanation: This is an example of the leveling effect. Water is a stronger acid than its conjugate base, OH−. Any base stronger than OH− will fully deprotonate water, so they all appear to have the same “strongest possible” basicity in aqueous solution.

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