Acid-base chemistry buffer solutions

Although this treatment of buffers was based on acid-base chemistry, the idea of a buffer is general and can be extended to equilibria involving complexation or redox reactions. For example, the Nernst equation for a solution containing Fe + and Fe + is similar in form to the Henderson-Hasselbalch equation. 

Buffer ions are used to maintain solutions at constant pH values. The selection of a buffer for use in the investigation of a biochemical process is of critical importance. Before the characteristics of a buffer system are discussed, we will review some concepts in acid-base chemistry. 

AQUEOUS SOLUTIONS AND ACID-BASE CHEMISTRY Weak Acids and Bases—Buffers... 

One of the most important applications of the acid-base chemistry is the buffer capacity. A buffer solution is a solution that counteracts an external action affecting pH. Some of the most well-known buffer solutions are found in the human body where they help to protect the pH of the blood from external actions and to keep the blood at constant pH levels. It is essential for the human body to be able to maintain the pH of the blood at fairly constant levels as certain types of cell only may survive in a close pH window. 

You learned about acids and bases in your previous chemistry course. In this chapter, you will extend your knowledge to learn how the structure of a compound determines whether it is an acid or a base. You will use the equilibrium constant of the reaction of an acid or base with water to determine whether the acid or base is strong or weak. You will apply your understanding of dissociation and pH to investigate buffer solutions solutions that resist changes in pH. Finally, you will examine acid-base titrations that involve combinations of strong and weak acids and bases. 

Potassium hydrogen phthalate has many uses in analytical chemistry. It is a primary standard for standardization of bases in aqueous solutions. Its equivalent weight is 204.2. It also is a primary standard for acids in anhydrous acetic acid. Other applications are as a buffer in pH determinations and as a reference standard for chemical oxygen demand (COD). The theoretical COD of a Img/L potassium hydrogen phthalate is 1.176mg O2. 

Equation 2.6 is the familiar Henderson-Hasselbalch equation, which defines the relationship between pH and the ratio of acid and conjugate base concentrations. The Henderson-Hasselbalch equation is of great value in buffer chemistry because it can be used to calculate the pH of a solution if the molar ratio of buffer ions ([A-]/[HA]) and the pKa of HA are known. Also, the molar ratio of HA to A- that is necessary to prepare a buffer solution at a specific pH can be calculated if the pKa is known. 

The term buffer is usually used to describe a buffered solution, which is a solution that resists changes to pH when acids or bases are added to it. You know enough chemistry right now to know that there are limits to the amount of acid or base that a solution can buffer. You can probably imagine that a 100 ml beaker of buffered solution won t be much good if you mix it with a liter of concentrated hydrochloric acid The amount of acid of base that a buffered solution can handle before appreciably changing pH is known as the buffer capacity. The purpose of this section is to review the nature of buffers and how this relates to the buffer capacity. 

June 3, 1978, Lynn, Massachusetts, USA - Feb. 10, 1942 Boston, USA) Henderson studied medicine at Harvard and was Professor of Biological Chemistry at Harvard University, Cambridge, Massachusetts, from 1904 to 1942 [i]. Henderson published on the physiological role of -> buffers [ii-vii] and the relation of medicine to fundamental science. Because he and also - Hasselbalch made use of the law of mass action to calculate the - pH of solutions containing corresponding acid-base pairs, the buffer equation is frequently (esp. in the biological sciences) referred to as -> Henderson-Hasselbalch equation. 

Very many problems in solution chemistry are solved with use of the acid and base equilibrium equations. The uses of these equations in discussing the titration of weak acids and bases, the hydrolysis of salts, and the properties of buffered solutions are illustrated in the following sections of this chapter. 

Control of pH is vital in synthetic and analytical chemistry, just as it is in living organisms. Procedures that work well at a pH of 5 may fail when the concentration of hydronium ion in the solution is raised tenfold to make the pH 4. Fortunately, it is possible to prepare buffer solutions that maintain the pH close to any desired value by the proper choice of a weak acid and the ratio of its concentration to that of its conjugate base. Let s see how to choose the best conjugate acid-base system and how to calculate the required acid-base ratio. 

Textbooks of analytical chemistry should be consulted for further details concerning the ionization of weak acids and bases and the theory of indicators, buffer solutions, and acid-alkali titrations. 

Further we introduced buffer chemistry and saw how pH may be calculated in buffer solutions and on how the buffer equation is often used in practice. When one has a solution of a weak acid and its corresponding weak base, both in concentrations of the same magnitude, one has a buffer system and the buffer equation may be used to calculate pH. Lastly, we looked at titration and on pH curves exemplifyed through examples of titration of monovalent weak acid with strong base and titration of divalent acid likewise with strong base NaOH. In the end we saw how colour indicator work. 

A buffer is something that resists change. In terms of acid and base chemistry, a buffer solution tends to resist change in pH when small to moderate amounts of a strong acid or strong base are added. A buffer solution consists of a mixture of a weak acid and its conjugate base. 

Solutions that contain a weak conjugate acid—base pair, such as those discussed in Section 17.1, resist drastic changes in pH when small amounts of strong acid or strong base are added to them. These solutions are called buffered solutions (or merely buffers). Human blood, for example, is a complex buffered solution that maintains the blood pH at about 7.4 (see the Chemistry and Life box on page 713). Much of the chemical behavior of seawater is determined by its pH, buffered at about 8.1 to 8.3 near the surface (see Chemistry and Life box on page 728). Buffered solutions find many important applications in the laboratory and in medicine ( FIGURE 17.1). 

The general chemistry of NAD(P)" and NAD(P)H will not be covered or further commented on in this chapter, except that NAD(P)H is relatively stable in aqueous solutions at pHs more alkaline than pH 7 and NAD(P)+ at pHs more acidic than pH 7. The stability is very much dependent on the buffer constituents and ionic strength. The pH where both redox forms together exhibit minimal destruction due to acid/base decomposition is found between pH 7 and 8, depending on whether the aqueous medium is unbuffered and when buffered, on the buffer constituents. In general, the stability of NADPH is less dependent on the buffer than is that of NADH. The reader is advised to refer to... 

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