Chapter 18: Aqueous Ionic Equilibria

Buffer Solutions: Definition, Components, and Properties

Buffer solutions are essential in maintaining a stable pH in various chemical and biological systems. They consist of specific combinations of weak acids and their conjugate bases, or weak bases and their conjugate acids, at predetermined concentrations or ratios.

• Definition: A buffer solution is a mixture that resists changes in pH when small to moderate amounts of strong acid or base are added, or when diluted.

• Types of Buffers:

  • Weak acid (HA) and its conjugate base (A–)

  • Weak base (B) and its conjugate acid (HB+)

• Role of Soluble Salts: Soluble salts are often used to supply the conjugate acid or base component. For example:

  • Acetic acid and sodium acetate (HC2H3O2/NaC2H3O2)

  • Ammonia and ammonium chloride (NH3/NH4Cl)

  • Carbonic acid and sodium bicarbonate (H2CO3/NaHCO3)

  • Sodium bicarbonate and sodium carbonate (NaHCO3/Na2CO3)

• Naturally Occurring Buffers:

  • Blood maintains a constant pH (~7.45) due to buffers such as the H2CO3/HCO3– system.

  • Natural water's resistance to acid rain is determined by its carbonic acid/bicarbonate buffer.

Understanding Buffer Action

The fundamental property of a buffer is its ability to resist pH changes upon addition of strong acids or bases, or upon dilution. This is governed by the equilibrium between the weak acid/base and its conjugate pair.

• Key Equilibrium: For a weak acid buffer:

• Buffer Response to Acid: Addition of strong acid increases [HA] and decreases [A–], but the pH changes minimally.

• Buffer Response to Base: Addition of strong base increases [A–] and decreases [HA], again with minimal pH change.

• Mechanism: The buffer components react with added H3O+ or OH– to maintain the equilibrium ratio, thus stabilizing pH.

Henderson-Hasselbalch Equation and pH Calculation

The Henderson-Hasselbalch equation provides a convenient way to calculate the pH of a buffer solution based on the concentrations of its components.

• Derivation: Generalized form:

• Application:

  • For weak acid/conjugate base: base = A–, acid = HA

  • For weak base/conjugate acid: base = B, acid = HB+

• Features:

  • If [base] = [acid], then pH = pKa

  • Each tenfold change in the ratio [base]/[acid] changes pH by 1 unit

  • Dilution does not affect pH, as the ratio remains constant

• Diprotic Acid Buffer Systems:

  • Two equilibria and two buffer systems:

• Example Calculations:

  • Acetic acid/sodium acetate buffer:

  • Ammonia/ammonium chloride buffer:

• Limitations: The Henderson-Hasselbalch equation is not valid when:

  • Acid dissociation constant (Ka) > 10–3

  • Molar concentration of acid or base < 10–3 M

  Additional info: In such cases, a quadratic equation must be used for accurate pH calculation.

Preparation of Buffer Solutions

Buffer solutions are prepared in the laboratory by mixing calculated amounts of acid and conjugate base (or base and conjugate acid) to achieve a desired pH and total concentration.

• Protocol:

  1. Obtain assigned pH and buffer components.

  2. Calculate the required amounts for 100.0 mL of buffer, ensuring total concentration equals 0.20 M.

  3. Use provided Ka values and molar masses for calculations.

• Example: For a sodium bioxalate/sodium oxalate buffer at pH 4.77:

  • Identify relevant acid dissociation equations (monoprotic, diprotic, triprotic).

  • Choose the correct acid/conjugate base pair and use the appropriate Ka value.

  • Calculate concentrations: Let , then Solve for : (HC2O4–), (C2O42–)

  • Calculate masses:

• General Preparation Tips:

  • Weigh solid components using an analytical balance.

  • If a component is a liquid, use its density to convert mass to volume and measure with a buret.

• pH Adjustment: Theoretical calculations may not match measured pH due to uncertainties in dissociation constants and simplifications. Adjust pH by dropwise addition of strong acid or base, monitoring with a pH meter and stirring continuously.

Buffer Capacity and Range

Buffer capacity refers to the amount of acid or base a buffer can neutralize before the pH changes significantly. The effective range of a buffer is typically within ±1 pH unit of its pKa.

• Key Points:

  • Buffer capacity increases with higher concentrations of buffer components.

  • Buffer range is determined by the pKa of the acid/base pair.

Calculations Involving Buffers Using the ICE Table

ICE (Initial, Change, Equilibrium) tables are used to track concentrations during buffer reactions, especially when calculating pH after addition of strong acid or base.

• Steps:

  1. List initial concentrations of buffer components.

  2. Calculate changes due to added acid/base.

  3. Determine equilibrium concentrations and use the Henderson-Hasselbalch equation for pH.

Table: Buffer System Examples

Additional info: Buffer systems are widely used in analytical chemistry, biochemistry, and environmental science to maintain pH stability.