A buffer is a solution that resists changes in pH upon the addition of acid or base.
Normally, the addition of acid to a solution will lower its pH and the addition of a base will raise its pH. If the solution is a buffer, however, its pH will be changed less than would be expected from the amounts of acid or base that are added.
Blood and many other bodily fluids are naturally buffered to resist changes in pH. Buffers are of great importance in living systems. Both the rates of biochemical reactions and their equilibrium constants are very sensitive to the availability of hydronium ions. Many biochemical reactions involved in vital processes like metabolism, respiration, the transmission of nerve impulses, and muscle contraction and relaxation take place within a narrow pH range. Le Cha^telier’s principle and the same types of chemical reactions that apply for the acetic acid-acetate buffer system govern the behavior of the physiological buffers.
An important buffer in the blood consists of bicarbonate ion and dissolved carbon dioxide in the form of carbonic acid. These two species constitute the conjugate acid-base pair of the buffering system. The pH of the blood can be altered by the ingestion of acidic or basic substances, and the carbonate/bicarbonate buffer system compensates for such additions and maintains the pH within the required range. This buffering system is intimately tied to respiration, and an exceptional feature of pH control by this system is the role of ordinary breathing in maintaining the pH.
Carbon dioxide is a normal product of metabolism. It is transported to the lungs, where it is eliminated from the body by exhalation. However, carbon dioxide in water—which is essentially equivalent to carbon dioxide in blood—hydrolyses to form carbonic acid, which dissociates to produce the hydrogen carbonate ion and the hydronium ion.
O2 (aq) + H2 O(l)← “H2 CO3” (aq)← H2 O+(aq) + HCO2-(aq)
If a chemical reaction or the ingestion of an acid increases the hydronium ion concentration in the blood, bicarbonate ions react with the added hydronium ions and are transformed into carbonic acid, which means that the concentration of dissolved carbon dioxide in the blood increases. Respiration increases, and more carbon dioxide is expelled from the lungs. In the terminology of Le Cha^telier’s principle, the conjugate acid-base equilibrium is stressed by the addition of acid. In response to that stress, the conjugate base reacts with the hydronium ion, producing more carbon dioxide, which is removed from the system by increasing respiration.
If a base is ingested, the hydronium ion reacts with it and causes a decrease in the concentration of hydronium ions. The equilibrium compensates for this change by shifting to the right, so that more carbonic acid dissociates to restore the hydronium ions consumed by the base. This requires more carbon dioxide to be dissolved in the blood, so respiration is decreased and more gas retained. In the language of Le Cha^telier, the base reacts with the hydronium ion, stressing the equilibrium for its production to the left to restore balance. More carbonic acid dissociates to compensate for the hydronium ion consumed; in turn, respiration decreases so more CO2 is retained in the lungs to restore the carbonic acid concentration.
Because the pH of the blood is in large part under respiratory control, simple alterations in normal breathing can change the pH of the blood. The state of respiratory acidosis arises as a result of hypoventilation. Slow, shallow breathing causes more CO2 to be retained in the lungs, which in turn causes more CO2 to be dissolved in water in lungs and blood, more carbonic acid to be formed, and the concentration of bicarbonate and hydronium ion in the blood to rise. All equilibria involved are stressed toward the production of more hydronium ion; hence, the pH drops. Some drugs are alkaline, and one of the clinical signs of alkaline drug overdose is hypoventilation. Hyperventilation induces the opposite situation. Increased expulsion of CO 2 from the lungs causes the equilibrium to be stressed to the left, consuming hydronium ion and resulting in a pH more basic than the normal 7.4.
Dissociate— To break up into ions. Various compounds dissociate to various degrees when dissolved in water.
Electrolyte— The chemical solution in which an electric current is carried by the movement and discharge of ions.
LeChâtelier’s principle— When a stress is applied to a system in equilibrium, the system shifts its balance (its amounts of reactants and products) in whichever direction partially relieves the stress.
pH— A measure of the acidity of a solution that is, of the concentration of H+ ions in the solution. The pH is the logarithm of the molarity of H+ ions, with the sign changed.
Weak acid (base)— An acid (base) that dissociates less than 100% into H+ (OH–) ions when dissolved in water. An acid (base) that is a weak electrolyte.
In the kidneys, the buffer system is more complicated and involves the dihydrogen phosphate/monohydrogen phosphate system in addition to the bicarbonate buffer of blood. If acidity increases, the hydrogen phosphate ion acts as a base and accepts the added protons, thereby forming more dihydrogen phosphate. If a base consumes the hydronium ions, the dihydrogen phosphate dissociates to restore the hydronium ion concentration, thereby forming more hydrogen phosphate. In either case, the added acid or base is consumed by its reaction with the appropriate component of the buffer and the pH does not change much.
To act as a good buffer, a solution must maintain a nearly constant pH when either acid or base is added. Two factors must be considered when a buffer is prepared. First, what pH must be maintained? The desired pH defines the range of the buffer. Second, how much acid or base does the solution need to consume without a significant change in pH? This defines the capacity of the buffer. The desired pH also determines the conjugate acid-base pair used in making up the buffer. The quantity of acid or base the buffer must be able to consume determines the concentrations of the buffers components that must be used.
One of the most important applications of the acid-base properties of salts is the formation of buffer solutions. The pH of a buffer solution changes to a relatively small extent when acid or base is added to it. A solution containing a weak acid and the anion that is its conjugate base or a solution containing a weak base and the cation that is its conjugate acid acts as a buffer. Two crucial properties of a buffer are its range, the portion of the pH scale over which a particular buffer is effective, and its capacity, the amount of acid or base that can be added without causing a large change in pH. Buffers in living systems maintain the pH in a range that prevents denaturation (destruction and decomposition) of proteins and degradation of other pH-sensitive biomolecules and which allows biological reactions to take place consistently.
Berg, Jermey M., John L. Tymoczko, and Lubert Stryer. Biochemistry. New York: W.H. Freeman, 2006.
Nelson, David L. and Michael M. Cox. Lehninger Principles of Biochemistry, Fourth Edition. New York: W.H. Freeman, 2004.
Voet, Donald and Judith G. Voet. Biochemistry. New York: John Wiley & Sons, 2006.
Robert L. Wolke
In chemistry , a buffer is a system, usually an aqueous (water ) solution , that resists having its pH changed when an acid or a base is added to it.
Normally, the addition of acid to a solution will lower its pH and the addition of a base will raise its pH. If the solution is a buffer, however, its pH will be changed to a much lesser extent than would be expected from the amounts of acid or base that are added. Socalled "buffered aspirin" is not really a buffer, because it does not resist acids and bases . It is simply aspirin combined with a basic compound, such as magnesium carbonate or aluminum hydroxide , which neutralizes some stomach acid.
Almost all chemical reactions that take place in aqueous solution—meaning almost all chemical reactions— are sensitive to the concentrations of hydrogen ions and hydroxide ions, that is, to the pH of the solution. This is because hydrogen and hydroxide ions are the ions of water itself. In particular, many biochemical processes essential to life are quite sensitive to the acidities of various body fluids. A variety of natural buffer systems keep the body's pH values within the limits that are necessary for health. For example, a system of several buffers holds the pH of human blood between 7.33 and 7.43 in a healthy person. A blood pH below 7.0 or above 7.8 can be fatal.
How buffers work
There are two common kinds of buffer solutions: solutions that contain a weak acid plus one of its salts (e.g., acetic acid plus sodium acetate) and solutions that contain a weak base plus one of its salts (e.g., ammonia plus ammonium chloride). Their workings can be understood in terms of LeChâtelier's principle.
Weak acid buffers
When a weak acid is dissolved in water, only a few of its molecules dissociate to form only a few hydrogen ions; the rest of the acid molecules remain as undissociated, neutral molecules that do not affect the pH. For example, whenever acetic acid is added to water, the following three species will be in the solution:
To make a buffer solution out of this system, we can add many more acetate ions to the solution in the form of sodium acetate, which is a strong electrolyte and dissociates completely. Ignoring the sodium ions that come along with the sodium acetate because they do not affect the acidity at all, we then have:
This solution will resist having its hydrogen ion concentration changed. To see how that works, first consider what would happen if we were to add some acid-some extra hydrogen ions-to this solution. According to LeChâtelier's principle, the equilibrium will be shifted to the left. That is, the added hydrogen ions will react with some of the acetate ions to form more acetic acid molecules. The result is that almost all of the added hydrogen ions are used up to form "harmless" neutral molecules; they therefore are not available to increase the acidity of the solution. The solution has resisted having its pH lowered more than a little bit.
What if we were to add some base-hydroxide ions to the buffer solution? Hydroxide ions to react with hydrogen ions, because the resulting molecule , H2O, is so stable.
Therefore, the added hydroxide ions will quickly remove hydrogen ions from the buffer solution, which according to LeChâtelier's principle will then shift its equilibrium to the right, making more acetate ions out of acetic acid molecules. Thus, the added hydroxide ions will have been used up, and only "harmless" acetate ions will have been formed. (Acetate ions are slightly basic, however, so the pH of the buffer solution does increase slightly.)
Weak base buffers
When ammonia gas is added to water, it forms a solution of a weak base whose equilibrium can be represented as follows:
To make a buffer solution out of this system, we can add many more ammonium ions to the solution in the form of ammonium chloride, which is a strong electrolyte and dissociates completely. Ignoring the chloride ions that come along with the ammonium chloride because they do not affect the acidity at all, we will then have:
If we add hydrogen ions to this buffer solution, they will be neutralized by reacting with the hydroxide ions to form water. According to LeChâtelier's principle, this removal of hydroxide ions will shift the equilibrium to the right, producing more ammonium ions, which does not affect the pH (but ammonium ions are slightly acidic). If we add hydroxide ions to the ammonia buffer solution, they will shift the equilibrium to the left, which uses up the added hydroxide ions and forms more whole NH3molecules, which does not affect the pH.
Some important buffers
Common weak acid buffer systems are based upon carbonic acid (H2CO3), citric acid (H3C6H5O), and phosphoric acid (H3PO4). Common weak base buffer systems are based upon amines (organic bases) or amino acids, which can act as both acids and bases.
A buffer solution based upon a dibasic or tribasic acid (an acid that can produce two or three hydrogen ions per molecule) may be made from two different ions of the same acid, rather than from the acid itself and one salt ion. For example, the phosphoric acid ions H2PO- 4 and HPO2- 4 can form what is known as the phosphate buffer system, which is one of the buffers that control the pH of human blood. In this system the H2PO – 4 ion plays the role of the weak acid and the HPO2- 4 ion plays the role of its salt. The relevant equilibrium is
The main buffer that is involved in controlling blood pH, however, is the carbonate system, which is based on the following equilibrium:
Our breathing (oxygen in, carbon dioxide out) controls the amount of carbon dioxide that is available to dissolve in the bloodstream. Therefore, our lungs also play an important part in controlling our blood's pH through the carbonate buffer system.
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Ouellette, Robert. Chemistry: An Introduction to General, Organic, and Biological Chemistry. Prentice Hall, 1994.
Robert L. Wolke
KEY TERMS. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
- Amino acid
—To break up into ions. Various compounds dissociate to various degrees when dissolved in water.
—The chemical solution in which an electric current is carried by the movement and discharge of ions.
- LeChâtelier's principle
—When a stress is applied to a system in equilibrium, the system shifts its balance (its amounts of reactants and products) in whichever direction partially relieves the stress.
—A measure of the acidity of a solutionthat is, of the concentration of H+ ions in the solution. The pH is the logarithm of the molarity of H+ ions, with the sign changed.
- Weak acid (base)
—An acid (base) that dissociates less than 100% into H+ (OH-) ions when dissolved in water. An acid (base) that is a weak electrolyte.
A buffer is a solution that resists changes in pH upon the addition of acid or base. Buffers typically contain several species that react with added acid and base.
Buffers are important in maintaining the proper environment within microorganisms and within other cells, including those in man. In the microbiology laboratory, many solutions and growth media are buffered to prevent sudden and adverse changes in the acidity or alkalinity of the environment surrounding the microorganisms.
Blood is an example of a natural buffer. In water, small volumes of an acid or base solution can greatly change the pH (measure of the hydrogen ion concentration). If the same amount of the acid or base solution is added to blood, the normal pH of the blood (7.4) changes only marginally. Blood and many other bodily fluids are naturally buffered to resist changes in pH.
In order to explain the properties of a buffer, it is useful to consider a specific example, the acetic acid/acetate buffer system. When acid (e.g., HCl, hydrochloric acid) is added to this buffer, the added hydronium ion (H +) reacts with the strongest base in the medium, namely the acetate ion, to form more acetic acid. This reaction uses up the added hydronium ion, preventing the pH from rising drastically, and is responsible for the buffering effect. As a result of adding acid to the buffer, the concentration of acetate decreases and the concentration of acetic acid increases. The solution acts as a buffer because nearly all of the added hydronium ion is consumed by reaction with acetate. As the hydrogen ion concentration increases, the acetate concentration and acetic acid concentration must adjust. The pH changes slightly to reflect the shift in the concentrations, but the change is much smaller than in the absence of the buffer because most of the added acid is consumed by its reaction with the acetate ion. This example of an acetic acid/acetate ion buffer is typical of other buffer systems.
Buffers are vitally important in living prokaryotic and eukaryotic systems. The rates of various biochemical reactions are very sensitive to the availability of hydronium ions. Many biochemical reactions (e.g., metabolism , respiration , the transmission of nerve impulses, and muscle contraction and relaxation) take place only within a narrow range of pH.
An important buffer in the blood is the bicarbonate ion and dissolved carbon dioxide in the form of carbonic acid. The acidity or alkalinity of the blood can be altered by the ingestion of acidic or basic substances. The carbonate/bicarbonate buffer system compensates for such additions and maintains the pH within the required range.
This buffering system is intimately tied to respiration, and an exceptional feature of pH control by this system is the role of ordinary breathing in maintaining the pH. Carbon dioxide is a normal product of metabolism. It is transported to the lungs, where it is eliminated from the body with every exhalation. However, carbon dioxide in blood is converted to carbonic acid, which dissociates to produce the hydrogen carbonate ion and the hydronium ion. If a chemical reaction or the ingestion of an acidic material increases the hydronium ion concentration in the blood, bicarbonate ion reacts with the added hydronium ion and is transformed into carbonic acid. As a result the concentration of dissolved carbon dioxide in the blood increases. Respiration increases, and more carbon dioxide is expelled from the lungs. Conversely, if a base is ingested, the hydronium ion reacts with it, causing a decrease in the concentration of hydronium ion. More carbonic acid dissociates to restore the hydronium ion consumed by the base. This requires more carbon dioxide to be dissolved in the blood, so respiration is decreased and more gas is retained.
To act as a buffer, a solution must maintain a nearly constant pH when either acid or base is added. Two considerations must be made when a buffer is prepared: (1) Which pH is desired to maintain? The desired pH defines the range of the buffer. (2) How much acid or base does the solution need to consume without a significant change in pH? This defines the capacity of the buffer. The desired pH also determines the compounds used in making up the buffer. The quantity of acid or base the buffer must be able to consume determines the concentrations of the components that must be used, and which allows biological reactions to take place consistently.
See also Biochemical analysis techniques; Laboratory techniques in microbiology
When base is added, most of the extra hydroxide ions are removed by reaction with undissociated acid: OH– + H2CO3 → HCO3– + H2O
Thus, the addition of acid or base changes the pH very little. Basic buffers have a weak base and a salt of the base (to provide the conjugate acid).Natural buffers occur in living organisms, where the biochemical reactions are very sensitive to change in pH (see acid–base balance). The main natural buffers are H2CO3/HCO3– and H2PO4–/HPO42– (see also haemoglobinic acid). Buffer solutions are also used in the laboratory (e.g. to keep microscopical preparations at their original pHs in order to prevent the formation of artefacts), in medicine (e.g. in intravenous injections), in agriculture, and in many industrial processes (e.g. fermentation processes).
1. A temporary memory for data, normally used to accommodate the difference in the rate at which two devices can handle data during a transfer. The buffer may be built into a peripheral device, such as a printer or disk drive, or may be part of the system's main memory. See buffering.
2. A means of maintaining a short but varying length of magnetic tape between the reels and the capstan and head area of a tape transport, in order that the acceleration of the tape at the reels need not be as great as that of the tape at the capstan. There are two principal types of buffer: tension arm and vacuum column. In the first, the tape passes over a series of rollers, alternate rollers being fixed in position and the rest being attached to a sprung pivoted arm, so that a variable length of tape is taken up in the resulting loops; in the second, the tape is drawn by a difference of pressure into a chamber whose width is just that of the tape. Vacuum column transports are more expensive and noisier but can handle higher tape speeds.
Streaming tape transports and many types of cartridge drives do not use buffers and are therefore limited to lower accelerations of the tape in the area of the head and (if there is one) capstan.
3. Any circuit or device that is put between two others to smooth changes in rate or level or allow asynchronous operation. For example, line drivers can be used to isolate (or buffer) two sets of data lines.
buff·er / ˈbəfər/ • n. 1. a person or thing that prevents incompatible or antagonistic people or things from coming into contact with or harming each other: family and friends can provide a buffer against stress.2. (also buffer solution) Chem. a solution that resists changes in pH when acid or alkali is added to it. Buffers typically involve a weak acid or alkali together with one of its salts.3. Comput. a temporary memory area or queue used when transferring data between devices or programs operating at different speeds.• v. [tr.] 1. lessen or moderate the impact of (something): the massage helped to buffer the strain.2. treat with a chemical buffer.
A term in environmental chemistry that refers to the capacity of a system to resist chemical change. Most often it is used in reference to the ability to resist change in pH . A system that is strongly pH buffered will undergo less change in pH with the addition of an acid or a base than a less well buffered system. A well buffered lake contains higher concentrations of bicarbonate ions that react with added acid. This kind of lake resists change in pH better than a poorly buffered lake. A highly buffered soil contains an abundance of ion exchange sites on clay minerals and organic matter that react with added acid to inhibit pH reduction.
See also Acid and base