4: Acid-Base Chemistry & the Henderson-Hasselbalch Equation

Acid-Base Chemistry and Henderson-Hasselbalch Equation

• Objectives:
  • Describe the relationship between pH and hydrogen ion concentration.
  • Relate pKa to the ratio of associated and dissociated compounds.
  • Estimate the percentage of a compound that is dissociated or associated given pKa and pH.
  • Describe the function of a buffer in a biological system.
  • Explain how the bicarbonate system is a physiological buffer.
  • Describe why a buffer is more effective near its pKa(s).
  • Compare the ability of a compound to cross a biological membrane in the charged vs. uncharged state.

What is pH?

• Hydrogen ion (H+) concentration in an aqueous solution.
• pH = -log[H+].
• Lower pH, higher H+ concentration.
• Logarithmic scale; a change of one pH unit corresponds to a tenfold change in H+ concentration.

Why is pH Important?

• pH affects the structure and activity of biological macromolecules.
• Small changes in pH can cause large changes in structure and function.
• Blood and urine pH are used in medical diagnoses.
• Blood pH < 7.4 = acidosis; Blood pH > 7.4 = alkalosis.
• The body has a buffering capacity to maintain pH.

Conjugate Acid-Base Pair

• pH is determined by H+.
• H+ depends on solutes functioning as acids (proton donors) or bases (proton acceptors).
• Biological systems contain weak acids and their conjugate bases.
• Strength of the tendency to lose or gain a proton in H2O.
• A proton donor and its corresponding proton acceptor form a conjugate acid-base pair.

Acid Dissociation Constant (Ka)

• Ka describes the affinity of an acid for the dissociable H+.
• Keq for the reaction is called the acid dissociation constant, Ka.
• Weak acids have a low Ka.
• High HA concentration.
• Strong acids have a high Ka.
• Most in dissociated form (A-).

Acid Dissociation Constant (Ka)

• Ka values are small.
• Ka is expressed as pKa.
• Weak acids have a high pKa.
• Strong acids have a low pKa.
• Examples include ammonia (pKa 9.25) and acetic acid (pKa 4.76).

Conjugate Acid-Base Pair (Examples)

• Monoprotic: Acetic acid (CH3COOH), Acetate (CH3COO-), Ammonium ion.
• Diprotic: Carbonic acid, Glycine.
• Triprotic: Phosphoric acid.

Titration Curve

• Titration is used to determine the amount of acid.
• Acid is titrated with a solution of a strong base (NaOH).
• Plot of pH against amount of NaOH added, revealing the pKa of the weak acid.

Titration Curve (Continued)

• As NaOH is added, OH- combines with H+ to form water (H2O).
• CH3COOH dissociates further, resulting in more acetate (CH3COO-).
• At midpoint, the concentration of the proton donor equals the concentration of the proton acceptor (pH = pKa).

Titration Curve (Additional Points)

• Three weak acids with different dissociation constants (acetic acid, dihydrogen phosphate, ammonium ion).
• Their highest Ka corresponds to their lowest pKa and strong acids have lower pKa.

Buffers

• Resist changes in pH when small amounts of acid (H+) or base (OH-) are added.
• Buffer systems consist of a weak acid and its conjugate base.
• Effective within ±1 pH unit of the pKa of the weak acid.

Acetic Acid-Acetate as a Buffer System

• Nearly equal concentrations of conjugate proton donor and acceptor.
• Small change in the ratio of weak acid and its anion results in small change in pH.

Henderson-Hasselbalch Equation

• pH = pKa + log ([A-]/[HA]).
• Relates pH, pKa, and the ratio of proton acceptor (A-) to proton donor (HA).
• Shows that the pKa of a weak acid is equal to the pH of the solution at the midpoint of its titration.

Henderson-Hasselbalch Equation (Continued)

• Knowing the pKa of an acid and the pH, one can predict if the acid is largely protonated or largely deprotonated.
• If pH > pKa, the ratio [A-]/[HA] > 1 and the acid is largely deprotonated.
• If pH < pKa, the ratio [A-]/[HA] < 1 and the acid is largely protonated.

Amino Acids Buffer Cells and Tissues

• Intracellular and extracellular fluids have a characteristic and nearly constant pH.
• Defense mechanism against changes in pH = buffer systems.
• Cytoplasm of most cells contains high concentrations of proteins.
• Functional groups (weak acids/bases) buffer effectively near neutral pH (e.g., histidine).

Ionization of Histidine in the Blood

• The pKa value of the imidazole side chain of histidine is 6.0.
• Blood pH is 7.3. The ratio of protonated to deprotonated histidine is estimated to be 10:1.

Phosphate Buffer System

• Dihydrogen phosphate (H2PO4-) acts as the proton donor.
• Hydrogen phosphate (HPO42-) acts as the proton acceptor.
• A buffer system at physiological pH, near neutral pH (approximately 6.86).
• Active in biological fluids, effective in extracellular and intracellular environments, at pH ranges 6.9–7.4.

Bicarbonate Buffer System

• Carbonic acid (H2CO3) is the proton donor.
• Bicarbonate (HCO3-) is the proton acceptor.
• CO2(aq) is in equilibrium with CO2(g).
• pH is ultimately determined by dissolved bicarbonate (HCO3-) and CO2 in biological fluids.

Bicarbonate Buffer System (Continued)

• Effective near pH 7.4.
• Involves three reversible reactions.
• CO2 in lungs and bicarbonate in the blood are involved in maintaining blood pH (controlled by brain stem).
• Detection of changes in CO2 and pH.

Bicarbonate Buffer System (Further Points)

• Hyperventilation upsets the balance of O2 and CO2, favoring CO2 removal, raising blood pH, and leading to alkalosis.
• Homes remedies for mild alkalosis include breathing into a paper bag, increasing CO2 in the blood.

Influence of pH on Ionizable Drugs (General Principles)

• Many drugs are weak acids that exist in solution as both nonionized/lipid-soluble and ionized/lipid-insoluble forms.
• Distribution across membranes is affected by pKa and pH of the environment.
• The pKa represents the pH at which half of the drug is in its ionized form.
• Higher pH → increased dissociation.
• Lower pH → decreased dissociation.

Dissociation of Aspirin

• Aspirin (acetylsalicylic acid) has a pKa of 3.5.
• At pH 3.5, aspirin is 1:1 associated:dissociated.
• Stomach pH is ~1.5.
• Aspirin in stomach is largely associated (protonated).
• Duodenum pH is ~ 6.5.
• Aspirin in duodenum is largely dissociated (deprotonated).
• A 10-fold change in H+ concentration corresponds to a one-unit change in pH.
• A 100- or 1000-fold change corresponds to two- or three-unit pH changes.

Estimating % Dissociated

• Given the pH and pKa, and knowing the ratio relationship of protonated to deprotonated forms, one can estimate the percentage dissociated.