Buffer Definition - Chemistry and Biology

Buffer Definition - 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 â‡Å' H   A− Examples of Buffers blood - contains a bicarbonate buffer systemTRIS bufferphosphate buffer As stated, buffers are useful over specific pH ranges. For example, here is the pH range of common buffering agents: Buffer pKa pH range citric acid 3.13., 4.76, 6.40 2.1 to 7.4 acetic acid 4.8 3.8 to 5.8 KH2PO4 7.2 6.2 to 8.2 borate 9.24 8.25 to 10.25 CHES 9.3 8.3 to 10.3 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. Its 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, McIvaines 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 bodys 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/ed038p559Hulanicki, 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. Vogels 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.