DPN-Biochem I - Lecture 3 - Ionic equilibria isoelectric point

Lecture Overview

  • Topic: Ionic equilibria

  • Key Concepts: Henderson–Hasselbalch equation, isoelectric point, electrophoresis

  • Context: Biochemistry (I)

Ionic Equilibria

  • Definition: Describes the equilibrium between ions in a solution.

  • Brønsted-Lowry Theory:

    • Acids: Proton donors.

    • Bases: Proton acceptors.

  • Strong Acids: Fully dissociate into hydronium ions and a stable conjugate base.

  • Weak Acids: Partially dissociate; reach equilibrium between dissociated and non-dissociated forms.

  • Strong Bases: Fully ionize, providing a high concentration of OH- ions.

  • Weak Bases: Participate in equilibria between protonated and deprotonated forms.

Water, Its Ions, & pH

  • Water Ionization: Water can form hydronium (H3O+) and hydroxide (OH-) ions through reversible ionization:

    • Equilibrium constant (Kw): [H+][OH-] = 1.0 x 10^-14 M @ 25°C.

    • Calculating pH: pH = -log[H+].

pH of NaOH Solution

  • A 0.1 M NaOH solution completely dissociates.

  • Calculation for [H+]:

    • [H+] = Kw/[OH-] = 1 x 10^-14 M / 0.1 M = 1 x 10^-13 M.

    • pH = - log (1 x 10^-13) = 13.

Weak Acids Equilibrium

  • General form: HA ⇌ H+ + A-

  • Equilibrium Constant (Keq):

    • Ka = [H+][A-] / [HA]

    • pKa = -log10(Ka)

  • Lower pKa indicates a stronger tendency for proton removal.

Examples of Weak Acids

Acid

Conjugate Base

pKa

Ka (M)

Formic Acid (HCOOH)

Formate Ion (HCOO-)

3.75

1.78 × 10^-4

Acetic Acid (CH3COOH)

Acetate Ion (CH3COO-)

4.76

1.74 × 10^-5

Phosphoric Acid (H3PO4)

Phosphate Ions (H2PO4-, HPO4^2-)

Various

Various

Henderson–Hasselbalch Equation

  • Formula: pH = pKa + log([A-]/[HA])

  • Describes pH changes during weak acid/base titration.

  • Useful for buffering capacity analysis.

Buffered Solutions

  • Function: Minimize pH changes by neutralizing added acid/base.

  • Best functioning near their pKa (pKa ± 1).

Common Buffers in Biochemistry

  • Including Acetate, Citrate, Phosphate, Carbonate with specified pH ranges and pKa values.

Carbonic Acid Buffer in Blood

  • CO2 + H2O ↔ H2CO3 ↔ HCO3- + H+ helps maintain blood pH around neutral.

  • pKa of carbonic acid: 6.3, provides buffering capacity in physiological pH range.

Isoelectric Point (pI)

  • Definition: The pH at which a molecule has no net charge.

  • Behavior of molecules in relation to pI:

    • pH < pI: Molecule is positively charged

    • pH > pI: Molecule is negatively charged

    • pH = pI: Molecule exists as a zwitterion (neutral).

Amino Acids and pKa

  • Amino acids are ampholytes (contain both acidic and basic groups).

  • Their pKa values depend on their structure and surrounding environment.

  • Titration can determine pI for amino acids and polypeptides.

Effect of pH on Molecular Charge

  • Changes in pH can shift the ratio of protonated ([HA]) vs. deprotonated ([A-]) forms.

Charges in Biomacromolecules

  • DNA carries a negative charge, affecting interactions with proteins.

  • Proteins can be positively charged at physiological pH due to basic residues.

Electrophoresis

  • Process: Movement of charged molecules in an electric field.

    • Negatively charged molecules move toward the anode; positives toward cathode.

  • Gel Electrophoresis: Utilizes porous gels (e.g., polyacrylamide, agarose) for size separation.

  • Capillary Electrophoresis: Requires smaller sample amounts and provides higher resolution.

  • Isoelectric Focusing: Separates proteins by charge, stopping at their pI.

Proteomics

  • Study of protein characteristics/functions in different conditions.

  • 2D electrophoresis used for enhanced resolution of protein expression.