Acid-Base Equilibria Quiz

Acid-base Equilibria

Definition

• Acid-base equilibria refers to the balance between the concentration of acids and bases in a solution
• Crucial in biological systems, as it affects the functioning of enzymes, proteins, and other biomolecules

pH Scale

• pH is a measure of the concentration of hydrogen ions (H+) in a solution
• pH = -log[H+]
• pH range: 0-14
• pH 7: neutral (neither acidic nor basic)
• pH < 7: acidic
• pH > 7: basic

Acid-Base Equilibrium Reactions

• Acid dissociation:
  • HA (acid) ⇌ H+ (proton) + A- (conjugate base)
  • Example: CH3COOH (acetic acid) ⇌ H+ + CH3COO-
• Base association:
  • B (base) + H+ (proton) ⇌ BH+ (conjugate acid)
  • Example: NH3 (ammonia) + H+ ⇌ NH4+

Henderson-Hasselbalch Equation

• Relates pH to the pKa of an acid and the concentrations of the acid and its conjugate base
• pH = pKa + log([A-]/[HA])
• pKa: negative logarithm of the acid dissociation constant (Ka)

Buffer Systems

• A buffer is a solution that resists changes in pH when an acid or base is added
• Composed of a weak acid and its conjugate base, or a weak base and its conjugate acid
• Buffering capacity: the ability of a buffer to resist pH changes
• Examples of buffer systems:
  • Bicarbonate buffering system (H2CO3 + HCO3-)
  • Phosphate buffering system (H2PO4- + HPO42-)

Importance of Buffer Systems

• Maintain a stable pH in biological systems
• Crucial for proper enzyme function and protein structure
• Help regulate the body's acid-base balance

Acid-base Equilibria

Definition

• Acid-base equilibria refers to the balance between the concentration of acids and bases in a solution
• Crucial in biological systems, as it affects the functioning of enzymes, proteins, and other biomolecules

pH Scale

• pH is a measure of the concentration of hydrogen ions (H+) in a solution
• pH = -log[H+]
• pH range: 0-14
• pH 7 is neutral, neither acidic nor basic
• pH < 7 is acidic
• pH > 7 is basic

Acid-Base Equilibrium Reactions

• Acid dissociation: HA ⇌ H+ + A-
• Example: CH3COOH (acetic acid) ⇌ H+ + CH3COO-
• Base association: B + H+ ⇌ BH+
• Example: NH3 (ammonia) + H+ ⇌ NH4+

Henderson-Hasselbalch Equation

• Relates pH to the pKa of an acid and the concentrations of the acid and its conjugate base
• pH = pKa + log([A-]/[HA])
• pKa is the negative logarithm of the acid dissociation constant (Ka)

Buffer Systems

• A buffer is a solution that resists changes in pH when an acid or base is added
• Composed of a weak acid and its conjugate base, or a weak base and its conjugate acid
• Buffering capacity is the ability of a buffer to resist pH changes
• Examples of buffer systems:
  • Bicarbonate buffering system (H2CO3 + HCO3-)
  • Phosphate buffering system (H2PO4- + HPO42-)

Importance of Buffer Systems

• Maintain a stable pH in biological systems
• Crucial for proper enzyme function and protein structure
• Help regulate the body's acid-base balance

Henderson-Hasselbalch Equation

• The Henderson-Hasselbalch equation describes the relationship between pH and pKa in a buffer system.

Equation

• pH = pKa + log10([A-] / [HA])
• The equation has four variables: pH, pKa, [A-], and [HA].

Variables

• pH: the negative logarithm of the hydrogen ion concentration
• pKa: the negative logarithm of the acid dissociation constant
• [A-]: the concentration of the conjugate base
• [HA]: the concentration of the acid

Key Concepts

• The pH of a buffer solution depends on the ratio of [A-] to [HA]
• When [A-] = [HA], the pH equals pKa, indicating the equivalence point
• If [A-] > [HA], the pH is greater than pKa, resulting in a basic solution
• If [A-] < [HA], the pH is less than pKa, resulting in an acidic solution

Applications

• The Henderson-Hasselbalch equation is used to calculate the pH of a buffer solution
• It helps determine the ratio of acid to conjugate base required for a specific pH
• Applications include biology, chemistry, and medicine, particularly in studying biological buffers and pH regulation