Buffer solution - wikipedia Buffer solution Acids and Bases Acid Acid -- base reaction Acid strength Acidity function Amphoterism Base Buffer solutions Dissociation constant Equilibrium chemistry Extraction Hammett acidity function pH Proton affinity Self - ionization of water Titration Acid types Brønsted Lewis Mineral Organic Strong Superacids Weak Monobasic acid Dibasic acid Tribasic acid Solid Base types Brønsted Lewis Organic Strong Superbases Non-nucleophilic Weak A buffer solution ( more precisely, pH buffer or hydrogen ion buffer ) is an aqueous solution consisting of a mixture of a weak acid and its conjugate base, or vice versa. Its pH changes very little when a small amount of strong acid or base is added to it. Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical applications. In nature, there are many systems that use buffering for pH regulation. For example, the bicarbonate buffering system is used to regulate the pH of blood. Contents 1 Principles of buffering 1.1 Buffer capacity 2 Applications 2.1 Simple buffering agents 2.2 `` Universal '' buffer mixtures 2.3 Common buffer compounds used in biology 3 Calculating buffer pH 3.1 Monoprotic acids 3.2 Polyprotic acids 4 See also 5 References 6 External links Principles of buffering Simulated titration of an acidified solution of a weak acid ( pK = 4.7 ) with alkali. Addition of hydroxide to an equilibrium mixture of a weak acid. HA, and its conjugate base, A Buffer solutions achieve their resistance to pH change because of the presence of an equilibrium between the acid HA and its conjugate base A. HA ⇌ H + A When some strong acid is added to an equilibrium mixture of the weak acid and its conjugate base, the equilibrium is shifted to the left, in accordance with Le Châtelier 's principle. Because of this, the hydrogen ion concentration increases by less than the amount expected for the quantity of strong acid added. Similarly, if strong alkali is added to the mixture the hydrogen ion concentration decreases by less than the amount expected for the quantity of alkali added. The effect is illustrated by the simulated titration of a weak acid with pK = 4.7. The relative concentration of undissociated acid is shown in blue and of its conjugate base in red. The pH changes relatively slowly in the buffer region, pH = pK ± 1, centered at pH = 4.7 where ( HA ) = ( A ). The hydrogen ion concentration decreases by less than the amount expected because most of the added hydroxide ion is consumed in the reaction OH + HA → H O + A and only a little is consumed in the neutralization reaction which results in an increase in pH. OH + H → H O Once the acid is more than 95 % deprotonated the pH rises rapidly because most of the added alkali is consumed in the neutralization reaction. Buffer capacity Buffer capacity, β, is a quantitative measure of the resistance of a buffer solution to pH change on addition of hydroxide ions. It can be defined as follows. β = d n d ( p ( H + ) ) ( displaystyle \beta = ( \frac ( dn ) ( d ( mathrm ( p ( H ^ ( + ) ) ) ) ) ) ) where dn is an infinitesimal amount of added base and d ( p ( H ) ) is the resulting infinitesimal change in the cologarithm of the hydrogen ion concentration. With this definition the buffer capacity of a weak acid, with a dissociation constant K, can be expressed as the d n d ( p H ) = 2.303 ( C A K a ( H + ) ( K a + ( H + ) ) 2 ) ( displaystyle ( \frac ( dn ) ( d ( mathrm ( pH ) ) ) ) = 2.303 left ( ( \frac ( C_ ( mathrm ( A ) ) K_ ( mathrm ( a ) ) ( mathrm ( H ^ ( + ) ) ) ) ( left ( K_ ( mathrm ( a ) ) + ( mathrm ( H ^ ( + ) ) ) \right ) ^ ( 2 ) ) ) \right ) ) Buffer capacity for a 0.1 M solution of an acid with pK of 7 where C is the analytical concentration of the acid. pH is defined as − log ( H ). This equation shows that there are three regions of raised buffer capacity. At very low p ( H ) the ( H ) term in the denominator predominates and buffer capacity rises exponentially with decreasing pH. The buffer capacity of a buffering agent is at a local maximum when p ( H ) = pK. It falls to 33 % of the maximum value at p ( H ) = pK ± 1 and to 10 % at p ( H ) = pK ± 1.5. For this reason the useful range is approximately pK ± 1. Buffer capacity is proportional to the concentration of the buffering agent, C, so dilute solutions have little buffer capacity. At very high p ( H ) the second term in the denominator is effectively constant and buffer capacity rises exponentially with increasing pH because of the ( H ) term in the numerator ; ( H ) = 10. Properties 1 and 3 are independent of the presence or absence of added buffering agents. They are concentration effects and reflect the fact that pH is related to the logarithm of the hydrogen ion concentration. Applications Buffer solutions are necessary to keep the correct pH for enzymes in many organisms to work. Many enzymes work only under very precise conditions ; if the pH moves outside of a narrow range, the enzymes slow or stop working and can denature. In many cases denaturation can permanently disable their catalytic activity. A buffer of carbonic acid ( H CO ) and bicarbonate ( HCO ) is present in blood plasma, maintaining a pH between 7.35 and 7.45. Industrially, buffer solutions are used in fermentation processes and in setting the correct conditions for dyes used in colouring fabrics. They are also used in chemical analysis and calibration of pH meters. The majority of biological samples that are used in research are made in buffers, especially phosphate buffered saline ( PBS ) at pH 7.4. Simple buffering agents Buffering agent pK Useful pH range Citric acid 3.13, 4.76, 6.40 2.1 -- 7.4 Acetic acid 4.8 3.8 -- 5.8 KH PO 7.2 6.2 -- 8.2 CHES 9.3 8.3 -- 10.3 Borate 9.24 8.25 -- 10.25 For buffers in acid regions, the pH may be adjusted to a desired value by adding a strong acid such as hydrochloric acid to the buffering agent. For alkaline buffers, a strong base such as sodium hydroxide may be added. Alternatively, a buffer mixture can be made from a mixture of an acid and its conjugate base. For example, an acetate buffer can be made from a mixture of acetic acid and sodium acetate. Similarly an alkaline buffer can be made from a mixture of the base and its conjugate acid. `` Universal '' buffer mixtures By combining substances with pK values differing by only two or less and adjusting the pH, a wide range of buffers can be obtained. Citric acid is a useful component of a buffer mixture because it has three pK values, separated by less than two. The buffer range can be extended by adding other buffering agents. The following mixtures ( McIlvaine 's buffer solutions ) have a buffer range of pH 3 to 8. 0.2 M Na HPO ( mL ) 0.1 M citric acid ( mL ) pH 20.55 79.45 3.0 38.55 61.45 4.0 51.50 48.50 5.0 63.15 36.85 6.0 82.35 17.65 7.0 97.25 2.75 8.0 A mixture containing citric acid, monopotassium phosphate, boric acid, and diethyl barbituric acid can be made to cover the pH range 2.6 to 12. Other universal buffers are the Carmody buffer and the Britton -- Robinson buffer, developed in 1931. Common buffer compounds used in biology For effective range see Buffer capacity, above. Common name Structure pK at 25 ° C Temp. effect dpH / dT ( K ) Mol. weight TAPS 8.43 − 0.018 243.3 Bicine 8.35 − 0.018 163.2 Tris 8.07 * − 0.028 121.14 Tricine 8.05 − 0.021 179.2 TAPSO 7.635 259.3 HEPES 7.48 − 0.014 238.3 TES 7.40 − 0.020 229.20 MOPS 7.20 − 0.015 209.3 PIPES 6.76 − 0.008 302.4 Cacodylate 6.27 138.0 MES 6.15 − 0.011 195.2 ( * ) Tris is a base, the pK of 8.07 refers to its conjugate acid. Calculating buffer pH Monoprotic acids First write down the equilibrium expression. HA ⇌ A + H This shows that when the acid dissociates equal amounts of hydrogen ion and anion are produced. The equilibrium concentrations of these three components can be calculated in an ICE table. ICE table for a monoprotic acid ( HA ) ( A ) ( H ) 0 y − x x x C − x x x + y The first row, labelled I, lists the initial conditions : the concentration of acid is C, initially undissociated, so the concentrations of A and H would be zero ; y is the initial concentration of added strong acid, such as hydrochloric acid. If strong alkali, such as sodium hydroxide, is added y will have a negative sign because alkali removes hydrogen ions from the solution. The second row, labelled C for change, specifies the changes that occur when the acid dissociates. The acid concentration decreases by an amount − x and the concentrations of A and H both increase by an amount + x. This follows from the equilibrium expression. The third row, labelled E for equilibrium concentrations, adds together the first two rows and shows the concentrations at equilibrium. To find x, use the formula for the equilibrium constant in terms of concentrations : K a = ( H + ) ( A − ) ( H A ) ( displaystyle K_ ( mathrm ( a ) ) = ( \frac ( ( mathrm ( H ^ ( + ) ) ) ( mathrm ( A ^ ( - ) ) ) ) ( ( mathrm ( HA ) ) ) ) ) Substitute the concentrations with the values found in the last row of the ICE table : K a = x ( x + y ) C 0 − x ( displaystyle K_ ( mathrm ( a ) ) = ( \frac ( x ( x + y ) ) ( C_ ( 0 ) - x ) ) ) Simplify to : x 2 + ( K a + y ) x − K a C 0 = 0 ( displaystyle x ^ ( 2 ) + ( K_ ( mathrm ( a ) ) + y ) x-K_ ( mathrm ( a ) ) C_ ( 0 ) = 0 ) With specific values for C, K and y this equation can be solved for x. Assuming that pH = − log ( H ) the pH can be calculated as pH = − log ( x + y ). Polyprotic acids % species formation calculated for a 10 millimolar solution of citric acid. Polyprotic acids are acids that can lose more than one proton. The constant for dissociation of the first proton may be denoted as K and the constants for dissociation of successive protons as K, etc. Citric acid, H A, is an example of a polyprotic acid as it can lose three protons. Equilibrium pK value H A ⇌ H A + H pK = 3.13 H A ⇌ HA + H pK = 4.76 HA ⇌ A + H pK = 6.40 When the difference between successive pK values is less than about three there is overlap between the pH range of existence of the species in equilibrium. The smaller the difference, the more the overlap. In the case of citric acid, the overlap is extensive and solutions of citric acid are buffered over the whole range of pH 2.5 to 7.5. Calculation of the pH with a polyprotic acid requires a speciation calculation to be performed. In the case of citric acid, this entails the solution of the two equations of mass balance C A = ( A 3 − ) + β 1 ( A 3 − ) ( H + ) + β 2 ( A 3 − ) ( H + ) 2 + β 3 ( A 3 − ) ( H + ) 3 C H = ( H + ) + β 1 ( A 3 − ) ( H + ) + 2 β 2 ( A 3 − ) ( H + ) 2 + 3 β 3 ( A 3 − ) ( H + ) 3 − K w ( H + ) − 1 ( displaystyle ( \begin ( aligned ) C_ ( ( ce ( A ) ) ) & = ( ( ce ( A ^ 3 - ) ) ) + \beta _ ( 1 ) ( ( ce ( A ^ 3 - ) ) ) ( ( ce ( H+ ) ) ) + \beta _ ( 2 ) ( ( ce ( A ^ 3 - ) ) ) ( ( ce ( H+ ) ) ) ^ ( 2 ) + \beta _ ( 3 ) ( ( ce ( A ^ 3 - ) ) ) ( ( ce ( H+ ) ) ) ^ ( 3 ) C_ ( ( ce ( H ) ) ) & = ( ( ce ( H+ ) ) ) + \beta _ ( 1 ) ( ( ce ( A ^ 3 - ) ) ) ( ( ce ( H+ ) ) ) + 2 \beta _ ( 2 ) ( ( ce ( A ^ 3 - ) ) ) ( ( ce ( H+ ) ) ) ^ ( 2 ) + 3 \beta _ ( 3 ) ( ( ce ( A ^ 3 - ) ) ) ( ( ce ( H+ ) ) ) ^ ( 3 ) - K_ ( ( ce ( w ) ) ) ( ( ce ( H+ ) ) ) ^ ( - 1 ) end ( aligned ) ) ) C is the analytical concentration of the acid, C is the analytical concentration of added hydrogen ions, β are the cumulative association constants log ⁡ β 1 = p K a 3, log ⁡ β 2 = p K a 2 + p K a 3, log ⁡ β 3 = p K a 1 + p K a 2 + p K a 3 ( displaystyle log \beta _ ( 1 ) = ( ce ( p ) ) K_ ( ( ce ( a3 ) ) ), quad log \beta _ ( 2 ) = ( ce ( p ) ) K_ ( ( ce ( a2 ) ) ) + ( ce ( p ) ) K_ ( ( ce ( a3 ) ) ), quad log \beta _ ( 3 ) = ( ce ( p ) ) K_ ( ( ce ( a1 ) ) ) + ( ce ( p ) ) K_ ( ( ce ( a2 ) ) ) + ( ce ( p ) ) K_ ( ( ce ( a3 ) ) ) ) K is the constant for self - ionization of water. There are two non-linear simultaneous equations in two unknown quantities ( A ) and ( H ). Many computer programs are available to do this calculation. The speciation diagram for citric acid was produced with the program HySS. See also Henderson -- Hasselbalch equation Buffering agent Good 's buffers Common - ion effect Metal ion buffer Mineral redox buffer References Jump up ^ Butler, J.N. ( 1964 ). Ionic Equilibrium : A Mathematical Approach. Addison - Wesley. p. 151. ^ Jump up to : Hulanicki, A. ( 1987 ). Reactions of acids and bases in analytical chemistry. Translated by Masson, Mary R. Horwood. ISBN 0 - 85312 - 330 - 6. Jump up ^ Scorpio, R. ( 2000 ). Fundamentals of Acids, Bases, Buffers & Their Application to Biochemical Systems. ISBN 0 - 7872 - 7374 - 0. Jump up ^ McIlvaine, T.C. ( 1921 ). `` A buffer solution for colorimetric comparaison '' ( PDF ). J. Biol. Chem. 49 ( 1 ) : 183 -- 186. Archived ( PDF ) from the original on 2015 - 02 - 26. Jump up ^ Mendham, J. ; Denny, R.C. ; Barnes, J.D. ; Thomas, M. ( 2000 ). `` Appendix 5 ''. Vogel 's textbook of quantitative chemical analysis ( 5th ed. ). Harlow : Pearson Education. ISBN 0 - 582 - 22628 - 7. Jump up ^ Carmody, Walter R. ( 1961 ). `` Easily prepared wide range buffer series ''. J. Chem. Educ. 38 ( 11 ) : 559 -- 560. Bibcode : 1961JChEd... 38... 559C. doi : 10.1021 / ed038p559. Jump up ^ `` Buffer Reference Center ''. Sigma - Aldrich. Archived from the original on 2009 - 04 - 17. Retrieved 2009 - 04 - 17. Jump up ^ Alderighi, L. ; Gans, P. ; Ienco, A. ; Peters, D. ; Sabatini, A. ; Vacca, A. ( 1999 ). `` Hyperquad simulation and speciation ( HySS ) : a utility program for the investigation of equilibria involving soluble and partially soluble species ''. Coordination Chemistry Reviews. 184 ( 1 ) : 311 -- 318. doi : 10.1016 / S0010 - 8545 ( 98 ) 00260 - 4. Archived from the original on 2007 - 07 - 04. External links `` Biological buffers ''. REACH Devices. Phosphate buffer Chemical equilibria Concepts Acid dissociation constant Binding constant Binding selectivity Buffer solution Chemical equilibrium Chemical stability Chelation Determination of equilibrium constants Dissociation constant Distribution coefficient Dynamic equilibrium Equilibrium chemistry Equilibrium constant Equilibrium unfolding Equilibrium stage Hammett acidity function Henry 's law Liquid -- liquid extraction Macrocycle effect Phase diagram Predominance diagram Phase rule Reaction quotient Self - ionization of water Solubility equilibrium Stability constants of complexes Thermodynamic equilibrium Thermodynamic activity Vapor -- liquid equilibrium Retrieved from `` https://en.wikipedia.org/w/index.php?title=Buffer_solution&oldid=854146253 '' Categories : Acid -- base chemistry Equilibrium chemistry Buffers Talk Contents About Wikipedia Български Bosanski Català Čeština Dansk Deutsch Eesti Ελληνικά Español Euskara فارسی Français Gaeilge Galego 한국어 हिन्दी Bahasa Indonesia Italiano עברית Қазақша Kreyòl ayisyen Latviešu Magyar Македонски മലയാളം Nederlands 日本 語 Norsk Polski Português Română Русский Simple English کوردی Српски / srpski Srpskohrvatski / српскохрватски Suomi Svenska தமிழ் Türkçe Українська Tiếng Việt 中文 35 more Edit links This page was last edited on 9 August 2018, at 08 : 07 ( UTC ). 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Question: what is the function of a buffer in chemistry
Answer: keeping pH at a nearly constant value in a wide variety of chemical applications