Buffer Definition and Examples in Chemistry


What Is a Buffer in Chemistry
A buffer solution resists a change in pH from the addition of a small amount of acid or base.

A buffer is a solution that maintains the stability of a system’s pH level when adding small quantities of acids or bases. This characteristic makes buffers important in biological and chemical applications where pH stability is crucial.

Composition of Buffer Solutions

A buffer solution typically consists of a weak acid and its conjugate base, or a weak base and its conjugate acid. These components work in tandem to neutralize any added acid or base.

Examples of Buffers

For example, the following acid and base pairs work together and form buffer solutions. A salt supplies the conjugate acid or base as it dissolves in solution.

  • The weak acid acetic acid (CH3COOH) and a salt containing the conjugate base (the acetate ion CH3COO), such as sodium acetate (CH3COONa)
  • The weak base ammonia (NH3) and a salt containing its conjugate acid (the ammonium cation NH4+), such as ammonium hydroxide (NH4OH)

How a Buffer Works

Buffers function through a process of chemical equilibrium. When you add an acid to a buffer, the conjugate base present in the buffer neutralizes it. Conversely, when you add a base, the weak acid in the buffer neutralizes the base.

For example, consider a buffer made of acetic acid (CH₃COOH) and sodium acetate (CH₃COONa). When hydrochloric acid (HCl) is added, the acetate ion (CH₃COO⁻) reacts with the H⁺ ions from HCl, forming more acetic acid and mitigating the pH change. Similarly, adding a base like sodium hydroxide (NaOH) results in the acetic acid reacting with OH⁻ to produce acetate and water, again stabilizing the pH.

There are limits to how much acid or base a buffer solution handles before it can no longer maintain the pH. Also, selecting the right buffer system depends on the desired pH.

Selecting the Best Buffer for a Desired pH

The choice of an appropriate buffer depends on the desired pH and the buffer’s pKa, the dissociation constant of the acid (or conjugate acid). A buffer is most effective when the pH is close to the pKa of the acid in the buffer system.

For example, the Ka of hydrofluoric acid is 6.6 x 10-4. Therefore its pKa value is -log(6.6 x 10-4) = 3.18. This means a hydrofluoric acid buffer works best around a pH of 3.18.

Meanwhile, the Kb value of the weak base ammonia (NH3) is 1.8 x 10-5. This means the Ka for its conjugate acid (NH4+) is Kw/Kb = 10-14 / 1.8×10-5 = 5.6 x 10-10. The pKa of NH4+ is -log(5.6×10-5) = 9.25. The optimal pH for the NH4+/NH3 buffer system is around a pH of 9.25.

Choose a weak acid and its salt for pH values that are lower than 7. Select a weak base and its salt for pH values that are above 7.

Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation estimates the pH of the solution when the amounts of conjugate acid and conjugate base are approximately equal (within a factor of 10). This equation relates pH, pKa, and the ratio of the concentration of the conjugate base to that of the weak acid:

pH = pKa + log([Base]/[Acid]​)

Rearrange the equation to calculate the required concentrations of the buffer components to achieve the desired pH.

Checking pH After Adding Strong Acids or Bases

After adding a strong acid or base to a buffer, the pH often shifts. The Henderson-Hasselbalch equation estimate the new pH, considering the change in the concentration of the acid/base pair. However, check the pH using a pH meter or indicators.

Universal Buffers

Universal buffers maintain a stable pH over a wide range. These are mixtures of several different buffers. A classic example is the Britton-Robinson buffer, which is a mixture of phosphoric acid, boric acid, and acetic acid. The Britton Robinson buffer maintains the pH over the range from 2 to 12. Another example of a universal buffer is citric acid, which has three pKa values. Universal buffers are especially useful in applications requiring a stable pH across a broad spectrum.

Biological Buffers

Biological buffers play a critical role in maintaining the pH balance in living organisms, ensuring the proper functioning of biological processes. Here are some notable examples:

  1. Bicarbonate Buffer System: One of the most important buffering systems in human blood, the bicarbonate buffer consists of carbonic acid (H₂CO₃) and its conjugate base, bicarbonate (HCO₃⁻). This buffer helps maintain the blood pH around 7.4, with carbonic acid acting as the weak acid and bicarbonate as the weak base.
  2. Phosphate Buffer System: Widely present in biological fluids, this buffer consists of dihydrogen phosphate (H₂PO₄⁻) and hydrogen phosphate (HPO₄²⁻). It plays a significant role in buffering pH changes in cells and tissues, and is particularly effective in the pH range of 6.8 to 7.4.
  3. Proteins as Buffers: Many proteins, including hemoglobin in red blood cells, act as effective buffers. They contain amino acids with acidic and basic side chains that donate or accept protons, helping to buffer changes in pH. Hemoglobin, for instance, buffers the blood by binding to hydrogen ions and carbon dioxide.
  4. Amino Acids: Free amino acids in cells also contribute to buffering. The amino group (–NH₂) accepts a proton, while the carboxyl group (–COOH) donates a proton, making amino acids capable of buffering in different pH ranges.
  5. Citrate Buffer: Citrate, a key intermediate in the citric acid cycle, also functions as a buffer in metabolic pathways. It buffers pH changes in the pH range of 3.0 to 6.2.
  6. Tris Buffer: Tris (tris(hydroxymethyl)aminomethane) is a common buffer in biochemical and molecular biology labs. It maintains a stable pH in various types of solutions, including cell culture media and buffer solutions for DNA/RNA extraction and PCR.
  7. HEPES Buffer: Widely used in cell culture, HEPES (4-(2-hydroxyethyl)-1-piperazineethanesulfonic acid) is a zwitterionic buffer that is effective in the pH range of 6.8 to 8.2. HEPES is popular for its minimal interference with biological processes.

References

  • Butler, J. N. (1998). Ionic Equilibrium: Solubility and pH Calculations. Wiley. ISBN 978-0-471-58526-8.
  • Carmody, Walter R. (1961). “Easily prepared wide range buffer series”. J. Chem. Educ. 38 (11): 559–560. doi:10.1021/ed038p559
  • Scorpio, R. (2000). Fundamentals of Acids, Bases, Buffers & Their Application to Biochemical Systems. ISBN 978-0-7872-7374-3.
  • Skoog, Douglas A.; West, Donald M.; Holler, F. James; Crouch, Stanley R. (2014). Fundamentals of Analytical Chemistry (9th ed.). Brooks/Cole. ISBN 978-0-495-55828-6.
  • Urbansky, Edward T.; Schock, Michael R. (2000). “Understanding, Deriving and Computing Buffer Capacity”. Journal of Chemical Education. 77 (12): 1640–1644. doi:10.1021/ed077p1640