What is buffer in chemistry

Thus, the [HF] is about 1 M and the [F - ] is close to 0. This will be especially true once we have added more F - , the addition of which will even further suppress the dissociation of HF.

Thus, [F - ] should be about 0. For mL of solution, then, we will want to add 0. Since we are adding NaF as our source of F - , and since NaF completely dissociates in water, we need 0. Thus, 0. Such dilute solutions are rarely used as buffers, however. Recall that the amount of F - in the solution is 0.

Let's double check the pH using the Henderson-Hasselbalch Approximation , but using moles instead of concentrations:. Now let's see what happens when we add a small amount of strong acid, such as HCl. The Cl - is the conjugate base of a strong acid so is inert and doesn't affect pH, and we can just ignore it. In fact, we already discussed what happens. The equation is:. 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.

Buffer capacity depends on the amounts of the weak acid and its conjugate base that are in a buffer mixture.

For example, 1 L of a solution that is 1. The first solution has more buffer capacity because it contains more acetic acid and acetate ion. Figure 4. The graph, an illustration of buffering action, shows change of pH as an increasing amount of a 0. Note the greatly diminished buffering action occurring after the buffer capacity has been reached, resulting in drastic rises in pH on adding more strong base. When an excess of hydrogen ion enters the blood stream, it is removed primarily by the reaction:.

The pH of human blood thus remains very near the value determined by the buffer pairs pKa, in this case, 7. Normal variations in blood pH are usually less than 0. This equation relates the pH, the ionization constant of a weak acid, and the concentrations of the weak acid and its salt in a buffered solution. Scientists often use this expression, called the Henderson-Hasselbalch equation , to calculate the pH of buffer solutions.

Lawrence Joseph Henderson — was an American physician, biochemist and physiologist, to name only a few of his many pursuits. 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. He eventually became a professor at Harvard and worked there his entire life.

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 wrote an equation in to describe the carbonic acid-carbonate buffer system in blood. 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 also founded the Fatigue Laboratory, at the Harvard Business School, which examined human physiology with specific focus on work in industry, exercise, and nutrition.

In , Karl Albert Hasselbalch — , a Danish physician and chemist, shared authorship in a paper with Christian Bohr in 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. The normal pH of human blood is about 7. The carbonate buffer system in the blood uses the following equilibrium reaction:.

The concentration of carbonic acid, H 2 CO 3 is approximately 0. Using the Henderson-Hasselbalch equation and the p K a of carbonic acid at body temperature, we can calculate the pH of blood:. Therefore, there must be a larger proportion of base than acid, so that the capacity of the buffer will not be exceeded. Lactic acid is produced in our muscles when we exercise. An enzyme then accelerates the breakdown of the excess carbonic acid to carbon dioxide and water, which can be eliminated by breathing.

In fact, in addition to the regulating effects of the carbonate buffering system on the pH of blood, the body uses breathing to regulate blood pH. Solutions that contain appreciable amounts of a weak conjugate acid-base pair are called buffers. A buffered solution will experience only slight changes in pH when small amounts of acid or base are added.

Addition of large amounts of acid or base can exceed the buffer capacity, consuming most of one conjugate partner and preventing further buffering action. 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:. 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.

In order to understand how a buffer works, consider the example of a buffer solution made by dissolving sodium acetate into acetic acid. The equation for the reaction is:. If a strong acid is added to this solution, the acetate ion neutralizes it:. Acid buffer solutions have a pH less than 7.

Commonly used acidic buffer solutions are a mixture of ethanoic acid and sodium ethanoate in solution, which have a pH of 4. You can change the pH of the buffer solution by changing the ratio of acid to salt, or by choosing a different acid and one of its salts. Alkaline buffer solutions have a pH greater than 7 and are made from a weak base and one of its salts.

A very commonly used example of an alkaline buffer solution is a mixture of ammonia and ammonium chloride solution.