CHEM1010 W2 L3

Introduction to Buffers and pH

  • Topic: Base buffers and pH, outlined in Chapter 17 of the textbook.

  • Explore properties, preparation, and calculations of buffer solutions.

Definitions and Importance of Buffers

  • Buffer Definition: A buffer solution contains a weak acid and its conjugate base (or a weak base and its conjugate acid).

  • Buffers are crucial in maintaining stable pH levels in biological and industrial processes.

  • Example: Buffer systems in blood chemistry regulate pH despite acid/base fluctuations.

Buffer Components

  • Weak Acid: Partially dissociates to produce H+ ions, establishing equilibrium.

  • Conjugate Base: Typically the salt of the weak acid that dissolves entirely, providing the base component of the buffer.

  • Example: Acetic acid (weak acid) and sodium acetate (conjugate base).

Buffer Preparation Methods

  1. Mixing Weak Acid with Its Salt: Example - acetic acid with sodium acetate.

  2. Adding Strong Acid/Base: Strong acid (HCl) to a weak base (NH3) will produce the conjugate acid (NH4+).

Buffer Functionality and Calculations

  • Buffers can neutralize small amounts of added acid or base without significant pH changes.

  • ICE Tables: Used to determine equilibrium states after acid/base additions.

    • Initial concentrations, changes due to reactions, and equilibrium concentrations must be recorded.

  • Henderson-Hasselbalch Equation: Relates pH to concentrations of a weak acid and its conjugate base, expressed as:

    • pH = pKa + log([A-]/[HA])

  • Understanding how the equation is derived from equilibrium expressions aids in solving buffer-related problems.

Impact of Buffer Concentrations on Effectiveness

  • Buffers are more effective when [weak acid] and [conjugate base] are similar.

  • Buffering Capacity: The amount of acid/base a buffer can neutralize without significant pH change.

  • Buffering Range: Effective pH range is typically ±1 pKa; choosing the right weak acid is essential.

Example Problem: Acetic Acid Buffer

  • Given: 0.1 M acetic acid and 0.11 M sodium acetate; pKa from tables (approximately 4.76).

  • Calculate pH using the Henderson-Hasselbalch equation:

    • pH = 4.76 + log(0.11/0.1) results in minor pH changes after addition of acids/bases.

    • Showing that changing ratios directly affects buffer capacity and pH.

Practical Considerations for Buffer Use

  • Effective Buffering: Approximately equal amounts of weak acid and conjugate base result in small pH changes when acids/bases are added.

  • Concentration matters: Higher concentrations generally increase buffering capacity; dilute buffers are less effective.

Summary of Buffer Properties

  • Buffers consist of a weak acid and conjugate base pair, resisting pH fluctuations.

  • The Henderson-Hasselbalch equation is key for calculations.

  • Understand the effects of adding strong acids/bases on buffer systems using ICE tables.

  • Evaluate buffering efficiency based on component ratios and concentrations.

Practical Buffer Calculation Example

  • Adding hydrochloric acid to an acetate buffer:

    • Determine initial concentrations, volume changes after acid addition, and resulting equilibrium concentrations.

    • Calculate final pH after the addition showing minor pH shift demonstrating buffer effectiveness.

Conclusion

  • Understanding buffers is essential for managing pH in various fields, including biology and chemistry.

  • Importance of constant assessment of solution contents throughout calculations to maintain accurate results.

  • Prepare for further practice and clarify concepts as needed.