Buffer Definition in Chemistry and Biology

A buffer is a solution containing either a weak acid and its salt or a weak base and its salt, which is resistant to changes in pH. In other words, a buffer is an aqueous solution of either a weak acid and its conjugate base or a weak base and its conjugate acid. A buffer may also be called a pH buffer, hydrogen ion buffer, or buffer solution.

Buffers are used to maintain a stable pH in a solution, as they can neutralize small quantities of additional acid of base. For a given buffer solution, there is a working pH range and a set amount of acid or base that can be neutralized before the pH will change. The amount of acid or base that can be added to a buffer before changing its pH is called its buffer capacity. 

The Henderson-Hasselbalch equation may be used to gauge the approximate pH of a buffer. In order to use the equation, the initial concentration or stoichiometric concentration is entered instead of the equilibrium concentration.

The general form of a buffer chemical reaction is:

HA ⇌ H⁺ + A⁻

Examples of Buffers

• blood - contains a bicarbonate buffer system
• TRIS buffer
• phosphate buffer

As stated, buffers are useful over specific pH ranges. For example, here is the pH range of common buffering agents:

When a buffer solution is prepared, the pH of the solution is adjusted to get it within the correct effective range. Typically a strong acid, such as hydrochloric acid (HCl) is added to lower the pH of acidic buffers. A strong base, such as sodium hydroxide solution (NaOH), is added to raise the pH of alkaline buffers.

How Buffers Work

In order to understand how a buffer works, consider the example of a buffer solution made by dissolving sodium acetate into acetic acid. Acetic acid is (as you can tell from the name) an acid: CH₃COOH, while the sodium acetate dissociates in solution to yield the conjugate base, acetate ions of CH₃COO^-. The equation for the reaction is:

CH₃COOH(aq) + OH^-(aq) ⇆ CH₃COO^-(aq) + H₂O(aq)

If a strong acid is added to this solution, the acetate ion neutralizes it:

CH₃COO^-(aq) + H⁺(aq) ⇆ CH₃COOH(aq)

This shifts the equilibrium of the initial buffer reaction, keeping the pH stable. A strong base, on the other hand, would react with the acetic acid.

Universal Buffers

Most buffers work over a relative narrow pH range. An exception is citric acid because it has three pKa values. When a compound has multiple pKa values, a larger pH range becomes available for a buffer. It's also possible to combine buffers, providing their pKa values are close (differing by 2 or less), and adjusting the pH with strong base or acid to reach the required range. For example, McIvaine's buffer is prepared by combining mixtures of Na₂PO₄ and citric acid. Depending on the ratio between the compounds, the buffer may be effective from pH 3.0 to 8.0. A mixture of citric acid, boric acid, monopotassium phosphate, and diethyl barbituic acid can cover the pH range from 2.6 to 12!

Buffer Key Takeaways

• A buffer is an aqueous solution used to keep the pH of a solution nearly constant.
• A buffer consists of a weak acid and its conjugate base or a weak base and its conjugate acid.
• Buffer capacity is the amount of acid or base that can be added before the pH of a buffer changes.
• An example of a buffer solution is bicarbonate in blood, which maintains the body's internal pH.

Sources

• Butler, J. N. (1964). Ionic Equilibrium: A Mathematical Approach. Addison-Wesley. p. 151.
• Carmody, Walter R. (1961). "Easily prepared wide range buffer series". J. Chem. Educ. 38 (11): 559–560. doi:10.1021/ed038p559
• Hulanicki, A. (1987). Reactions of acids and bases in analytical chemistry. Translated by Masson, Mary R. Horwood. ISBN 0-85312-330-6.
• Mendham, J.; Denny, R. C.; Barnes, J. D.; Thomas, M. (2000). "Appendix 5". Vogel's Textbook of Quantitative Chemical Analysis (5th ed.). Harlow: Pearson Education. ISBN 0-582-22628-7.
• Scorpio, R. (2000). Fundamentals of Acids, Bases, Buffers & Their Application to Biochemical Systems. ISBN 0-7872-7374-0.