Notes on Acids, Bases, and Equilibria in Organic Chemistry

Acid Strength, Equilibria, and Carboxylate Behavior in Organic Chemistry

• Organic acids (carboxylic acids) are relatively weak acids by principle: they have a low tendency to dissociate in water.

  • In water, the dissociation is limited; the denominator (undissociated acid, HA) remains large for many organic acids.
  • Example: acetic acid (in salad dressing or white vinegar) shows a large amount of undissociated HA in solution, leading to a small Ka value.

• Acid strength is quantified by the acid dissociation constant Ka, defined by the expression Ka = \frac{[H^+][A^-]}{[HA]}

  • The numerator represents the dissociated form (proton and conjugate base), the denominator the undissociated acid.
  • For organic acids, typical orders of magnitude observed are K_a \in {10^{-2}, 10^{-5}, 10^{-10}} (these illustrate strong to very weak relative tendencies to dissociate).

• The notion of acid strength is context-dependent: in organic chemistry, the strength is relative to the reaction partner (the base) present in a given reaction.

  • An acid’s ability to donate a proton depends on whether the reaction partner is a suitable base to accept it.
  • Thus, acid strength is not an absolute attribute; it is conditional on the conjugate base available in the system.

• The pKa scale converts the power of ten in Ka to a logarithmic measure:
  pKa = -\log{10} K_a

  • The pKa expresses the exponent in the Ka value; a higher pKa means a weaker acid.
  • Different acids differ in pKa by multiples of 1; a difference of 5 pKa units corresponds to a factor of 10^5 in K_a.
  • Example: if one acid has pKa = 5 and another has pKa = 0, then their Ka differ by a factor of 10^5.

• The transcript notes a hypothetical consideration: how can Ka be as large as 1? It can be 1 if the masses-action expression has equal numerator and denominator, i.e., Ka = 1 \;\Rightarrow\; [H^+][A^-] = [HA]

  • In this situation, the forward and reverse dissociation processes occur with equal propensity.

• Entropy and enthalpy in chemical processes:

  • Entropy tends to favor greater disorder, i.e., products with more degrees of freedom.
  • Restoring order (e.g., after dissolution) requires energy input; energy must be supplied to push systems back toward more ordered states (e.g., to reform bonds).
  • Enthalpy describes bond formation and stabilization; favorable enthalpy (formation of strong bonds) can drive reactions forward.
  • Some reactions are extremely favorable due to a combination of enthalpy and entropy changes; in extreme cases this can lead to explosions, where stable molecules (e.g., CO2, CO, N2O, N_2) are formed along with energy release.
  • Conversely, some processes (e.g., photosynthesis) are not spontaneous under ambient conditions and require energy input (the sun) to overcome unfavorable ΔG and drive the reaction uphill to store energy as chemical bonds in sugars.

• Equilibrium and spontaneity:

  • At equilibrium, the free energy change is zero: \Delta G = 0.
  • A reaction will favor the side with the weaker acid (larger pKa) when comparing acid-base couples; a large difference in pKa (greater than about 5 units) strongly biases the equilibrium toward the weaker acid side.
  • If the pKa difference is > 5 in the direction that favors the left side, the equilibrium lies essentially on the left (the left-hand species is the weaker acid).
  • A smaller or more modest difference yields a gray zone where neither side is fully dominant; both forward and reverse processes have appreciable contributions.

• Practical implications: solubility and driving reactions with carboxylates

  • Carboxylate salts are ionic and generally highly water-soluble; ions dissolve readily in water.
  • By shifting an equilibrium toward the carboxylate salt (e.g., R-COO^- with Na^+), you can make species soluble in water that were previously insoluble in their protonated form.
  • A common lab strategy is to use a strong base (e.g., NaOH) to deprotonate a carboxylic acid, converting it to the carboxylate anion (the conjugate base) and a counterion (e.g., Na^+). This increases solubility and changes basicity.

• Relationship between acidity and basicity (conceptual)

  • The strength of a conjugate base is inversely related to the strength of its conjugate acid.
  • If one acid is stronger (lower pKa) than another, its conjugate base is weaker than the other acid's conjugate base.
  • In practice, choosing a base for driving a reaction is guided by the relative pKa values of the acids involved.

• Structural principles underlying acidity

  • Proton dissociation is influenced not only by bond polarization but also by bond length:
  • Longer, weaker bonds dissociate more readily, contributing to higher acidity.
  • Shorter, stronger bonds resist dissociation.
  • Thus, acidity is a function of both bond characteristics and the stability of the conjugate base.

• Important caveat about labeling

  • Simply calling something a “strong acid” or a “strong base” is not meaningful in isolation.
  • The strength label must be paired with the specific reaction context and partners involved (i.e., which reaction is taking place and which species are present).

• Quick self-check prompts (as in the transcript)

  • Any questions?

• Quick numerical recap from the transcript for reference:

  • Common orders of magnitude for organic acid dissociation constants: K_a \in {10^{-2},\;10^{-5},\;10^{-10}}
  • A difference of five pKa units corresponds to a factor of 10^5 in K_a.
  • A pKa difference of ten units (e.g., 5 vs 15) corresponds to a factor of 10^{10} in K_a.
  • In the gray zone around moderate pKa differences, practical manipulation (e.g., using bases like NaOH) can shift equilibria to form more soluble carboxylate salts.