Therefore, there must be a larger proportion of base than acid, so that the capacity of the buffer will not be exceeded. The fact that the H 2CO 3 concentration is significantly lower than that of the HCO 3 в€’ HCO 3 в€’ ion may seem unusual, but this imbalance is due to the fact that most of the by-products of our metabolism that enter our bloodstream are acidic. In 1916, Hasselbalch expressed HendersonвЂ™s equation in logarithmic terms, consistent with the logarithmic scale of pH, and thus the Henderson-Hasselbalch equation was born. The pH scale was introduced in 1909 by another Dane, SГёrensen, and in 1912, Hasselbalch published measurements of the pH of blood. In 1916, Karl Albert Hasselbalch (1874вЂ“1962), a Danish physician and chemist, shared authorship in a paper with Christian Bohr in 1904 that described the Bohr effect, which showed that the ability of hemoglobin in the blood to bind with oxygen was inversely related to the acidity of the blood and the concentration of carbon dioxide. He also founded the Fatigue Laboratory, at the Harvard Business School, which examined human physiology with specific focus on work in industry, exercise, and nutrition. Henderson was broadly knowledgeable in addition to his important research on the physiology of blood, he also wrote on the adaptations of organisms and their fit with their environments, on sociology and on university education. He wrote an equation in 1908 to describe the carbonic acid-carbonate buffer system in blood.

He discovered that the acid-base balance in human blood is regulated by a buffer system formed by the dissolved carbon dioxide in blood. He eventually became a professor at Harvard and worked there his entire life. He obtained a medical degree from Harvard and then spent 2 years studying in Strasbourg, then a part of Germany, before returning to take a lecturer position at Harvard. Lawrence Joseph Henderson (1878вЂ“1942) was an American physician, biochemist and physiologist, to name only a few of his many pursuits. Lawrence Joseph Henderson and Karl Albert Hasselbalch A change of 1 pH unit occurs when the acetic acid concentration is reduced to 11% of the acetate ion concentration. Figure 14.17 shows how pH changes for an acetic acid-acetate ion buffer as base is added. A buffer solution has generally lost its usefulness when one component of the buffer pair is less than about 10% of the other. A good buffer mixture should have about equal concentrations of both of its components.

There are two useful rules of thumb for selecting buffer mixtures: The first solution has more buffer capacity because it contains more acetic acid and acetate ion. For example, 1 L of a solution that is 1.0 M in acetic acid and 1.0 M in sodium acetate has a greater buffer capacity than 1 L of a solution that is 0.10 M in acetic acid and 0.10 M in sodium acetate even though both solutions have the same pH. Buffer capacity depends on the amounts of the weak acid and its conjugate base that are in a buffer mixture. The buffer capacity is the amount of acid or base that can be added to a given volume of a buffer solution before the pH changes significantly, usually by one unit. (credit: modification of work by Mark Ott) However, a large amount of acid exhausts the buffering capacity of the solution and the pH changes dramatically (beaker on the right). For example, strong base added to this solution will neutralize hydronium ion, causing the acetic acid ionization equilibrium to shift to the right and generate additional amounts of the weak conjugate base (acetate ion):ДЇigure 14.16 The indicator color (methyl orange) shows that a small amount of acid added to a buffered solution of pH 8 (beaker on the left) has little affect on the buffered system (middle beaker). The presence of a weak conjugate acid-base pair in the solution imparts the ability to neutralize modest amounts of added strong acid or base. To illustrate the function of a buffer solution, consider a mixture of roughly equal amounts of acetic acid and sodium acetate. (credit: modification of work by Mark Ott) How Buffers Work

(b) After the addition of 1 mL of a 0.01- M HCl solution, the buffered solution has not detectably changed its pH but the unbuffered solution has become acidic, as indicated by the change in color of the methyl orange, which turns red at a pH of about 4.

Figure 14.14 (a) The unbuffered solution on the left and the buffered solution on the right have the same pH (pH 8) they are basic, showing the yellow color of the indicator methyl orange at this pH.

Citect l buffer is to small