Biological amines play crucial roles in our bodies, but their behavior depends on pH. The Henderson-Hasselbalch equation helps us understand how these molecules exist in different forms at various pH levels, especially in our cells.
Knowing the protonation state of amines is key to grasping their function. This equation lets us calculate the ratio of neutral to protonated forms, giving insight into how these molecules behave in our bodies' slightly basic environment.
Biological Amines
Protonation states using Henderson-Hasselbalch equation
- Relates pH, pKa, and concentrations of weak acid and conjugate base: $pH = pKa + log([A^-]/[HA])$
- For weak bases, use pKb and ratio of base to conjugate acid: $pOH = pKb + log([B]/[BH^+])$
- Calculate protonation state at physiological pH (7.4):
- Find pKb of amine
- Convert pH to pOH: $pOH = 14 - pH$
- Plug pOH and pKb into rearranged equation
- Solve for $[B]/[BH^+]$ ratio
- $[B]/[BH^+]$ ratio shows relative amounts of neutral and protonated amine at given pH
- This equation is crucial for understanding buffer solutions in biological systems
Protonated form of cellular amines
- Cellular amines (amino acids, biogenic amines) often have pKb > physiological pH (7.4)
- When pKb > pH, amine more likely protonated
- Amine is stronger base than $OH^-$ at physiological pH
- Protonated form more stable under physiological conditions
- Writing in protonated form reflects predominant cellular state
- This concept is related to the Brønsted-Lowry theory of acids and bases
Neutral vs protonated weak bases
- Use Henderson-Hasselbalch for weak bases to find $[B]/[BH^+]$ at given pH
- Total amine concentration ($C_T$) = neutral ($[B]$) + protonated ($[BH^+]$):
- $C_T = [B] + [BH^+]$
- Let $x = [B]$, then $[BH^+] = C_T - x$
- Substitute into equation and solve for $x$
- Neutral percentage: $%B = (x / C_T) \times 100%$
- Protonated percentage: $%BH^+ = ((C_T - x) / C_T) \times 100%$
Henderson-Hasselbalch Equation
Protonation states using Henderson-Hasselbalch equation
- Relates pH, pKa, and concentrations of weak acid and conjugate base: $pH = pKa + log([A^-]/[HA])$
- For weak bases, use pKb and ratio of base to conjugate acid: $pOH = pKb + log([B]/[BH^+])$
- Calculate protonation state at physiological pH (7.4):
- Find pKb of amine
- Convert pH to pOH: $pOH = 14 - pH$
- Plug pOH and pKb into rearranged equation
- Solve for $[B]/[BH^+]$ ratio
- $[B]/[BH^+]$ ratio shows relative amounts of neutral and protonated amine at given pH
Neutral vs protonated weak bases
- Use Henderson-Hasselbalch for weak bases to find $[B]/[BH^+]$ at given pH
- Total amine concentration ($C_T$) = neutral ($[B]$) + protonated ($[BH^+]$):
- $C_T = [B] + [BH^+]$
- Let $x = [B]$, then $[BH^+] = C_T - x$
- Substitute into equation and solve for $x$
- Neutral percentage: $%B = (x / C_T) \times 100%$
- Protonated percentage: $%BH^+ = ((C_T - x) / C_T) \times 100%$
Acid-Base Equilibria and Dissociation Constants
- The Henderson-Hasselbalch equation is based on the concept of ionic equilibrium
- Acid dissociation constant (Ka) and base dissociation constant (Kb) are key parameters in understanding acid-base behavior
- pKa is the negative logarithm of Ka, used in the Henderson-Hasselbalch equation for acids
- pKb is the negative logarithm of Kb, used in the Henderson-Hasselbalch equation for bases